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LABOEATOEY   PEAOTIOE. 


LABORATORY    PRACTICE 


A  SERIES  OF  EXPERIMENTS  ON  THE 
FUNDAMENTAL   PRINCIPLES   OF    CHEMISTRY 


A    COMPANION   VOLUME   TO 

THE  NEW  CHEMISTRY 


BY 


JOSIAH  PAESOJSS  COOKE,  LL.  D. 

•sf- 

ERVING  PROFESSOR  AND  DIRECTOR  OF  THE  CHEMICAL  LABORATORY, 
HARVARD  UNIVERSITY 


NEW    YORK 
D.    APPLETON    AND    COMPANY 

1906 


COPYRIGHT,  1891, 
BY  JOSIAH  P.  COOKE. 


TO   MY    KINSMAN, 

FRANCIS    BARTLETT,   ESQ., 

WHOSE   SYMPATHY   AND   LIBERALITY 

HAS   ENCOURAGED   AND    PROMOTED   THE    STUDY 

OF   CHEMISTRY. 


CONTENTS. 


CHAPTER  PAGE 

INTRODUCTION 5 

I.  DISTINGUISHING  PROPERTIES 13 

1.  Water 14 

2.  Air  as  an  Example  of  Aeriform  Matter.        .        .  43 

3.  Oxygen  Gas 53 

4.  Hydrogen  Gas 59 

5.  Sulphur       .        . 64 

6.  Chlorine 70 

7.  Carbon .        .        .73 

8.  Nitrogen 81 

9.  Magnesium  and  Zinc 89 

10.  Sodium .        .92 

11.  Copper 97 

12.  Iron     .        .        . 100 

II.  GENERAL  PRINCIPLES 107 

13.  Province  of  Chemistry 107 

14.  Fundamental  Laws 110 

15.  Compounds  and  Elements 113 

16.  Qualitative  Analysis 116 

17.  Quantitative  Analysis 121 

III.  MOLECULES  AND  ATOMS 126 

18.  Molecular  Theory 126 

19.  Physical  Method  of  determining  Molecular  Weights  127 

20.  Chemical  Method  of  determining  Molecular  Weights  135 


4  LABORATORY  PRACTICE. 

CHAPTER  PAGE 

21.  Conception  of  Atoms 139 

22.  Determination  of  Atomic  Weights  .                .        .  144 
IV.  SYMBOLS  AND  NOMENCLATURE 149 

23.  Chemical  Symbols 149 

24.  Chemical  Reactions 154 

25.  Stochiometry 160 

26.  Nomenclature 162 

V.  MOLECULAR  STRUCTURE 164 

27.  Quantivalence 164 

VI.  THERMAL  RELATIONS 178 

28.  Heat  of  Chemical  Action.  .       .       .178 


INTRODUCTION. 


THIS  book  is  not  intended  to  be  nsed  withont  a 
teacher.  As  far  as  possible,  directions  are  given 
which  will  enable  the  student  to  perform  the  ex- 
periments successfully ;  but  in  many  cases  more 
precise  directions  and  cautions  are  required,  which 
should  be  given  in  the  lecture  room.  Whenever 
possible  the  experiments  should  be  performed  by 
the  teacher  in  sight  of  the  class  before  they  are 
attempted  by  the  students,  and  an  example  of  the 
necessary  apparatus  should  be  placed  on  the  lect- 
ure table  or  in  the  laboratory  where  it  can  be  in- 
spected. The  student  ought  to  be  left  to  make 
his  own  observations,  and  then  to  interpret  the 
results  with  such  aid  as  may  be  necessary  from 
the  instructor,  who  should  always  be  present  in 
the  laboratory  to  oversee  the  work.  The  educa- 
tional value  of  such  a  course  as  is  here  outlined 
depends  entirely  on  the  manner  in  which  the  work 
is  directed  and  supervised.  The  student  should 
be  instructed,  by  continued  reiteration,  if  neces- 
sary— 


6  LABORATORY  PRACTICE. 

1.  To  observe  the  minutest  particular  in  regard 
to  every  experiment. 

2.  To  distinguish  essential  from  non-essential 
phenomena. 

3.  To  draw  correct  inferences  from  the  results. 

4.  To  express  concisely  but  clearly  in  writing 
the  facts  observed  and  conclusions  reached. 

While  the  student  is  at  work  the  note  book 
should  be  always  on  the  laboratory  desk,  and  tak- 
ing notes  of  the  experiments  at  the  time  should  be 
most  strictly  enforced.  Remember  that  the  notes 
are  of  no  value  except  as  original  records  of  ob- 
servations, and  all  copying  of  notes  should  be  as 
absolutely  forbidden  as  is  the  tampering  with 
original  entries  in  mercantile  accounts.  If  cor- 
rections are  required  they  should  plainly  appear 
as  such. 

The  notes  should  be  so  written  that  an  exam- 
iner, not  necessarily  the  teacher,  can  gain  from 
them  in  the  shortest  possible  time  a  full  concep- 
tion of  the  work  done;  and  this  result  is  best 
secured  by  the  use  of  prominent  headings  and 
paragraphs  to  indicate  every  transition  in  the 
subject  matter.  Remember  that  an  examiner  can 
not  be  expected  to  decipher  hieroglyphics,  and 
that  his  impression  of  the  work  will  depend  in  no 
small  measure  on  the  clearness  with  which  the  re- 
sults are  presented. 

All  ciphering  required  in  order  to  reduce  the 


INTRODUCTION.  7 

results  of  experiments  or  to  answer  the  ques- 
tions in  this  book  should  be  made  in  the  note 
book  and  the  purport  of  every  calculation  clearly 
indicated  by  headings.  Experience  shows  that 
mere  arithmetical  mistakes  are  the  most  frequent 
sources  of  error  in  experimental  work,  and  if,  as 
is  the  too  frequent  practice,  the  separate  pieces  of 
paper  on  which  the  calculations  are  made  are 
thrown  away,  the  clue  to  the  error  is  often  lost. 
Of  course,  neatness  in  the  note  book  is  highly  de- 
sirable ;  but  this  may  be  secured  if  all  the  odd 
folios  are  reserved  for  notes  and  the  even  folios 
for  ciphering  and  the  headings  are  sufficiently 
multiplied  to  prevent  confusion. 

It  is  not  best  to  give  too  precise  directions  in 
regard  to  the  note  book,  but  to  allow  sufficient 
liberty  to  encourage  originality  both  in  form  and 
in  material,  otherwise  one  note  book  of  a  class 
of  students  will  be  a  precise  transcript  of  every 
other.  The  following  suggestions,  however,  will 
be  valuable  guides  to  a  good  result.  In  reference 
to  every  experiment — 

1.  State  the  materials  taken.* 

*  In  preparing  the  original  work  the  author  was  indebted  for 
the  above  rules  and  for  other  valuable  suggestions  to  an  excellent 
pamphlet  entitled  Laboratory  Experiments  in  General  Chemistry, 
by  William  Ripley  Nichols  and  Lewis  M.  Norton,  for  the  Use  of 
Students  of  the  Massachusetts  Institute  of  Technology.  Printed, 
not  published,  Boston,  1884.  In  rewriting  the  work  for  the  present 
edition  the  author  has  further  profited  by  the  experience  of  the  as- 


8  LABORATORY  PRACTICE. 

2.  Give  a  brief  description,  with  a  clear  sketch 
of  the  apparatus  employed. 

3.  State  the  method  used,  and  all  the  circum- 
stances of  the  experiment ;  that  is,  whether  a  sub- 
stance is  heated  or  cooled,  whether  a  gentle  or 
strong  heat  is  necessary,  whether  an  acid  employed 
is  strong  or  dilute,  etc. 

4.  Describe  what  takes  place ;  that  is,  all  that 
is  observed  during  and  as  a  result  of  the  experi- 
ment. 

5.  Give  the  conclusions  drawn,  or  state  what 
the  experiment  teaches. 

The  student  should  be  given  to  understand 
clearly  that  experiments  performed  mechanically, 
without  intelligence,  or  carelessly  recorded,  are 
worth  absolutely  nothing,  and  should  be  so  esti- 
mated in  any  system  of  school  or  college  credits. 

Experiments  are  only  of  value  as  parts  of  a 
course  of  instruction  logically  followed  out  from 
beginning  to  end.  In  such  a  course  there  must  be 
necessarily  a  great  deal  to  be  filled  out  by  the 
teacher ;  and  this  can  vastly  better  be  taught  from 
his  lips,  with  such  illustrations  as  he  can  command, 
than  from  any  books.  The  author  has  therefore 
thought  it  necessary  to  draw  up  in  connection 
with  the  experiments  an  outline  of  such  a  course 

sistants  who  have  had  charge  of  the  course  designated  at  Harvard 
College  as  Chemistry  B,  and  has  received  from  them  valuable  sug- 
gestions on  many  points. 


INTRODUCTION.  9 

as  lie  deems  best  suited  to  prepare  students  for 
the  further  study  of  natural  science  in  Harvard 
College,  stating  the  chief  points  to  be  illustrated, 
and  giving  the  order  in  which  they  are  best  pre- 
sented, leaving  it  to  the  teacher  to  fill  out  the  de- 
tails from  his  own  knowledge.  In  rewriting  the 
work  he  has  also  added  a  large  number  of  ques- 
tions and  problems  in  order  to  direct  attention  to 
important  inferences  which  might  otherwise  be 
overlooked  or  to  enforce  principles  which  are 
wont  to  be  imperfectly  apprehended  by  element- 
ary students.  It  is  intended,  however,  by  this  syn- 
opsis only  to  offer  suggestions  to  the  teacher,  and 
to  define  more  clearly  the  college  requisition  in 
chemistry  as  one  of  the  " Advanced  Studies"  in 
the  new  scheme  of  examinations  for  admission. 

It  is  thought  by  the  writer  that  a  course  on  the 
fundamental  principles  of  chemistry,  like  the  one 
here  outlined,  is  far  more  suitable  for  the  pupils 
of  secondary  schools  than  the  meagre  description 
of  the  scheme  of  the  chemical  elements  which  is 
presented  in  epitome  by  most  of  the  elementary 
text  books  on  this  science  ;  and  in  order  to  bring 
the  experimental  method  within  the  means  of  all 
schools  of  that  class,  the  writer  has  sought  to 
adapt  to  the  purposes  of  instruction  common 
household  utensils,  such  as  may  be  made  by  a 
tinsmith  or  found  at  any  house-furnishing  store. 
The  small  petroleum  or  gas  cooking  stoves  serve 


10  LABORATORY  PRACTICE. 

an  admirable  purpose  for  heating,  and  their  ovens 
are  excellent  drying  chambers  ;  a  farina  kettle  is 
a  good  steam  bath  ;  and  the  quick-sealing  fruit  or 
milk  jars  are  not  only  good  gas  holders,  but  ena- 
ble any  student  to  perform  experiments  which  for- 
merly were  made  only  with  costly  apparatus. 
Gflass  flasks,  retorts,  beaker  glasses,  glass  tubing 
and  tube  apparatus,  as  also  filtering  paper,  rubber 
tubing,  rubber  stoppers  and  corks,  besides  the 
chemicals  required,  are  best  purchased  of  regular 
dealers  in  chemical  supplies.  From  these  dealers 
may  also  be  procured  various  simple  devices  for 
supporting  apparatus,  although  equally  serviceable 
tools  can,  if  necessity  requires,  be  made  with  blocks 
of  wood  and  stout  iron  wire.  The  only  apparatus 
of  precision  required,  the  scales  and  thermome- 
ters, can  be  imported  from  Germany  at  the  prices 
hereafter  named  ;  and  the  whole  cost  of  an  outfit 
for  a  class  of  ten  students  should  not  exceed  two 
hundred  dollars,  and  much  less  will  be  required  if 
the  teacher  uses  a  little  ingenuity  and  can  spend 
time  in  extemporizing  apparatus  from  materials 
on  hand.  But  the  best  apparatus  will  be  of  no 
use  unless  the  teacher  stands  before  it  and  speaks 
to  his  pupils  out  of  the  fulness  of  his  own  knowl- 
edge. This  is  an  essential  condition  of  success, 
and  without  it  the  experimental  method  should 
never  be  attempted. 

In  devising  a  course  of  experiments  it  is  impos- 


INTRODUCTION.  H 

sible  to  foresee  every  contingency  ;  and  the  teach- 
er into  whose  hands  this  book  may  come  must  re- 
gard the  directions  here  given  simply  as  sugges- 
tions to  be  worked  out  experimentally  with  such 
appliances  as  he  can  command.  In  many  cases 
the  directions  can  undoubtedly  be  improved,  and 
no  experiment  should  ever  be  intrusted  to  an  in- 
experienced student  which  has  not  first  been  thor- 
oughly tested  by  the  teacher  ;  and  the  writer  is 
not  responsible  for  failures  which  must  result 
from  a  disregard  of  this  essential  precaution. 

NEWPORT,  August,  1891. 


DESCRIPTIVE  LIST  OF  CHEMICAL 
EXPERIMENTS 

INTENDED  TO   ILLUSTRATE  THE   GENERAL  PRINCIPLES 
OF  THE  SCIENCE. 


CHAPTER  I. 
DISTINGUISHING    PROPERTIES.* 

MATERIAL  bodies  are  made  up  of  an  unend- 
ing variety  of  substances.  Chemical  science  deals 
with  the  relations  of  these  substances.  In  a  world 
consisting  of  only  one  substance  the  relations  of 
mass  and  energy,  which  are  studied  in  physics, 
might  be  observed,  but  there  could  be  no  chem- 
istry. The  distinctions  of  substance  are,  there- 

*  Under  the  above  heading,  Distinguishing  Properties,  it  is 
expected  that  the  student  should  acquire  much  of  the  informa- 
tion usually  presented  in  elementary  books  on  chemistry  in  connec- 
tion with  the  scheme  of  the  chemical  elements,  but  the  facts  should 
be  studied  in  this  connection  simply  as  definite  phenomena  within 
the  range  of  personal  observation.  Many  teachers,  mistaking  the 
scheme  of  the  book,  not  probably  sufficiently  elaborated  in  the  first 
two  editions,  have  practised  their  pupils  from  the  first  in  writing  the 
reactions  of  the  processes  employed.  But  this  practice  not  only 
diverts  attention  from  the  main  purpose  of  an  experimental  course, 
but  so  confounds  fact  and  theory  in  the  mind  of  the  student  that  it 
becomes  subsequently  very  difficult  to  clear  up  the  confusion.  It 
must  not  be  forgotten  that  chemical  reactions  can  not  be  intelli- 
gently written  until  the  main  features,  at  least,  of  the  theory  of 
chemistry  have  been  fully  apprehended;  and  that  to  use  symbols 
wechanically  or  to  learn  formula  by  rote  stultifies  the  whole  study 
2 


14  LABORATORY   PRACTICE. 

fore,  the  basis  of  our  science.  In  most  cases 
these  distinctions  are  so  striking  as  to  be  obvious. 
No  person  would  confound  one  with  another 
salt,  sugar,  water,  or  air.  But  in  many  cases  the 
distinctions  are  obscure,  and  can  only  be  made 
manifest  by  careful  observation  and  study.  Hence 
a  philosophical  treatment  of  our  subject  implies  a 
preliminary  consideration  of  the  distinguishing 
properties  of  substances,  and  this  discussion  will 
give  the  opportunity  of  illustrating  many  funda- 
mental facts  and  relations  which  are  not  only  of 
great  interest  in  themselves,  but  which  will  also 
serve  as  the  basis  of  a  knowledge  of  the  principles 
of  chemical  science. 

1.  Water. 

EXPERIMENT  1.  Density  of  Water. — A  closed 
cylinder,  about  5  centimetres  long  by  2£  centi- 
metres in  diameter,  of  sheet  tin  or  brass,  with  a 
hook  at  the  top,  which  can  be  made  by  a  tin- 
smith, is  required  for  this  experiment.  Before 
soldering  on  the  top  the  cylinder  should  be 
loaded  with  lead  shot,  so  that  it  will  sink  in 
water.  A  pair  of  scales,  with  a  set  of  weights, 

of  physical  science.  To  make  evident  the  effect  of  ordinary  text- 
book teaching  the  following  question  was  asked  for  a  series  of  years 
in  the  examination  papers  on  chemistry  for  admission  to  Harvard 
College :  "  On  what  evidence  is  our  knowledge  of  the  composition 
of  water  based  ? "  and  in  at  least  four  cases  out  of  five  the  answer 
given  was :  "  Water  consists  of  hydrogen  and  oxygen,  because  H»  -j- 
O  =  HaO." 


DENSITY  OF  WATER.  15 

is  also  required.  Hand  scales,  with  horn  pans 
and  brass  beam  20  centimetres  long,  and  with 
a  set  of  weights  from  0*01  to  100  grammes,  are 
well  adapted  for  all  the  experiments  here  de- 
scribed, and  cost  in  Germany  only  $2.37.  The 
student  first  measures  the  size  of  the  cylinder. 
This  may  be  done  by  fitting  exactly  a  sheet  of 
glazed  paper  to  the  convex  surface,  and  measur- 
ing with  a  millimetre  scale  the  length  and  width 
of  the  paper.  The  volume  of  the  cylinder,  in 
cubic  centimetres,  can  now  be  calculated,  and  the 
student  should  be  shown  how  to  estimate  ap- 
proximately the  probable  error  of  his  result.  The 
cylinder  is  next  to  be  weighed,  first  in  air,  and 
secondly  under  iced  water.  For  this  purpose  the 
scales  are  best  hung  by  a  cord  passing  over  a  pul- 
ley and  secured  to  a  belaying-pin,  so  that  they 
can  be  adjusted  at  any  height  required.  The 
cylinder  is  best  suspended  beneath  the  pan  by  a 
silk  thread,  which  plays  freely  through  a  hole 
made  for  the  purpose  at  the  centre  of  the  pan, 
and  is  hung  from  the  same  hook  as  the  pan  itself. 
The  difference  between  the  weight  of  the  cylinder 
in  air  and  its  weight  under  water  will  now  give 
the  weight  of  a  volume  of  water  equal  to  that  of 
the  cylinder,  and  from  this  value  the  weight  of 
one  cubic  centimetre  of  ice-cold  water  is  easily 
found.  Since,  by  definition,  a  gramme  is  the 
weight  of  one  cubic  centimetre  of  pure  water  at  4° 


16  LABORATORY  PRACTICE. 

C.  (the  same  as  that  of  ice-cold  water  within  the 
limits  of  accuracy  of  student's  work),  the  result 
should  be  closely  one  gramme.  The  probable 
error  of  the  work  may  now  be  estimated  by  com- 
paring the  several  results  of  as  large  a  number  of 
different  students  as  possible.  Such  a  compari- 
son may  be  made  very  instructive  by  writing  the 
results,  distinguished  by  numbers  or  otherwise, 
in  a  vertical  column  on  the  blackboard,  and, 
after  finding  the  average  value,  placing  oppo- 
site to  each  result  the  difference  between  it  and 
this  average  value.  It  will  now  probably  appear 
that  some  of  the  results  are  far  astray,  in  conse- 
quence of  careless  work  or  mistakes  in  calcula- 
tion. These  should  be  thrown  out,  a  new  average 
value  taken,  and  a  still  closer  scrutiny  applied, 
when,  on  arranging  the  remaining  results  in  the 
order  of  their  values,  each  will  be  found  to  dif- 
fer from  the  next  by  a  small  and  nearly  constant 
quantity.  The  final  average  must  represent  very 
closely  the  best  result  that  can  be  obtained  with 
the  imperfect  means  employed ;  and  if  it  differs 
by  more  than  a  few  milligrammes  from  one 
gramme,  there  must  be  some  constant  source  of 
error  which  the  teacher  should  seek  to  discover. 
Such  a  comparison  as  this  not  only  furnishes  an 
unimpeachable  test  of  the  relative  skill  of  the  dif- 
ferent men  in  a  class  of  experimenters,  but  also 
gives  a  clearer  idea  of  the  necessary  limitations  of 


DENSITY  OF  WATER.  17 

experimental  methods  than  can  be  acquired  in 
any  other  way,  and  it  is  all-important  that  stu- 
dents should  gain  a  clear  idea  on  this  point  from 
the  start. 

NOTES,  QUESTIONS,  AND  PROBLEMS* 

(1)  What  is  the  principle  of  Archimedes,  and  by  what 
simple  consideration  can  you  show  that  it  must  be  true  ? 

(2)  How  can  you  apply  this  principle  to  determine  the 
specific  gravity  of  a  solid  body  ? 

(3)  In  this  book  by  specific  gravity  is  always  meant  the 
ratio  between  the  weight  of  a  substance  and  that  of  an  equal 
volume  of  water  or  that  of  some  other  standard  material, 
and  by  density  is  always  to  be  understood  the  weight  of  the 
unit  volume  of  a  substance.     In  the  French  system  the 
density  of  a  substance  is  the  weight  of  a  cubic  centimetre  of 
the  material  in  grammes.     In  the  English  system  it  is  the 
weight  of  a  cubic  inch  of  the  material  in  grains.     Specific 
gravity,  then,  is  a  ratio,  but  density  is  an  absolute  weight. 
If  we  multiply  the  density  of  the  standard  of  reference  (that 
is,  the  weight  of  the  unit  of  volume)  by  the  specific  gravity 
of  a  substance  we  shall  have  the  density  of  that  substance — 
that  is,  S  =  8'  Sp.  Gr.     Further,  if  we  have  given  the  volume 
of  a  body  and  the  specific  gravity  of  the  material  of  which 
it  consists,  the  weight  of  such  body  must  be  equal  to  the 
product  of  the  weight  of  one  unit  of  volume  (its  density)  into 
the  number  of  units  of  volume,  W  =  S  V  =  d'  Sp.  Gr.  V,  in 
which  8'  is  the  density  of  standard  of  reference  as  above.   In 
the  French  system,  since  by  definition  the  gramme  is  the 
weight  of  one  cubic  centimetre  of  water,  it  is  true  when  the 
standard  of  reference  is  water  that 

W  =  Sp.  Gr.  V, 

W  standing  for  a  certain  number  of  grammes  and  V  for  a 
certain  number  of  cubic  centimetres.  With  any  other  sys- 

*  This  heading  will  not  be  repeated,  but  must  be  regarded  as 
applying  to  the  numbered  paragraphs  in  small  type,  which  will 
very  constantly  follow  the  descriptions  of  experiments. 


18  LABORATORY   PRACTICE. 

tern  of  weights  and  measures  we  must  retain  in  the  formula 
the  value  8'  when 

W  =  «'  Sp.  Gr.  V. 

In  the  English  system  W  stands  for  weight  in  grains,  V  for 
volume  in  cubic  inches,  and  S'  for  the  weight  of  one  cubic 

fc 

inch  of  water  in  grains.    In  all  systems  Sp.  Gr.  =  — .    In  the 

French  system  alone  where  8'  =  1  gramme  we  have  Sp.  Gr. 
=  8.  This  equality  obviously  arises  solely  from  the  selec- 
tion of  a  cubic  centimetre  of  water  as  the  unit  of  weight  • 
but  while  this  assumption  greatly  simplifies  many  calcula- 
tions, and  is  one  of  the  chief  merits  of  the  French  system  of 
weights  and  measures,  it  has  led  to  a  confusion  of  ideas  as 
well  as  of  terms  which  it  is  important  to  keep  distinct. 

(4)  In  the  original  scheme  for  the  French  system  of  meas- 
ures and  weights,  the  metre  was  denned  as  the  one  ten-mill- 
ionth part  of  the  quadrant  of  a  meridian  of  the  earth, 
and,  in  order  to  construct  a  bar  corresponding  to  that  defini- 
tion, French  engineers  actually  measured  an  arc  of  the 
meridian  which  passes  through  Dunkirk,  in  northern 
France,  and  Barcelona,  in  Spain,  determining  with  extreme 
care  by  astronomical  means  the  latitude  of  these  two  places, 
so  as  to  find  what  ratio  the  arc  measured  bore  to  the  quad- 
rant. As  in  all  geodetic  measurements  on  a  large  scale,  the 
end  points  were  connected  by  a  system  of  triangulation,  re- 
ferred to  certain  base  lines  actually  measured  at  convenient 
localities  by  means  of  bars  of  standard  length.  These  bars 
were  necessarily  graduated  to  the  measures  of  length  pre- 
viously in  use  in  France,  and  as  the  result  of  the  work  the 
length  of  the  quadrant  was  found  (approximately)  in  terms 
of  the  old  standard  called  a  toise,  and  then  there  was  only 
the  mechanical  difficulty  left  of  making  a  metallic  bar  equal 
to  a  certain  fraction  of  a  toise,  and  this  was  the  new  standard 
since  called  the  metre.  A  metre  rule  having  been  thus 
constructed  with  the  subdivisions  accurately  marked  upon  it, 
the  weight  of  one  cubic  centimetre  of  water  (the  new  unit  of 
weight)  was  then  found  by  a  method  in  all  respects  similar 
to  that  employed  in  the  above  experiment.  A  metallic  cyl- 
inder was  made  much  larger  than  those  here  used,  and  care- 


EXPANSION  OP  WATER  BY  HEAT.  19 

fully  wrought  in  a  lathe  so  as  to  be  as  nearly  as  possible  of 
uniform  dimensions.  These  dimensions  were  next  meas- 
ured with  the  metric  scale  with  extreme  accuracy  and  the  vol- 
ume of  the  cylinder  in  cubic  centimetres  calculated.  It  was 
then  only  necessary  to  weigh  the  cylinder  in  the  air  and 
under  water,  when  the  chief  data  were  obtained  for  calculat- 
ing the  weight  of  one  cubic  centimetre  of  water  the  required 
gramme.  We  say  the  chief  data  because  all  the  weights 
and  measurements  had  to  be  corrected  for  variations  in 
temperature  or  in  other  conditions  which  it  would  be  out 
of  place  to  discuss  here.  But  the  result  of  all  was  to  give 
the  weight  of  the  new  standard  in  terms  of  the  old  sys- 
tem of  French  weights  used  in  the  work.  In  few  words, 
it  was  found  that  a  gramme  equalled  so  many  old  French 
grains,  and  afterwards  standard  gramme  weights  were 
made  in  brass  or  platinum  to  correspond  to  the  value  thus 
found. 

In  order  that  the  student  may  fully  understand  these 
somewhat  complex  relations,  he  should  repeat  the  above  ex- 
periment with  English  weights,  and  thus  find  the  weight  of 
a  cubic  centimetre  of  water  in  English  grains,  when  he  can 
himself  adjust  a  gramme  weight  with  a  piece  of  sheet  lead. 
If  this  is  not  possible  from  the  want  of  an  accurate  set  of 
English  weights,  he  may  be  set  a  problem  in  this  form  : 
Given  a  cylinder  of  such  dimensions  (in  centimetres),  weigh- 
ing so  many  grains  in  air  and  so  many  under  water,  what  is 
the  value  of  a  gramme  in  grains  ? 

1  gramme  =  15*432  grains. 

Ex.  2.  Expansion  of  Water  by  Heat.— Pro- 
vide a  cylindrical  glass  bulb  about  20  milli- 
metres wide  and  30  millimetres  long  opening  into 
a  tube  about  2  millimetres  wide  and  200  milli- 
metres long.  In  order  to  facilitate  the  empty- 
ing and  drying  of  the  bulb,  a  short  tubulature 
of  the  same  size  as  the  stem  should  be  provided 
at  the  opposite  end  and  drawn  out  so  that  it 


20  LABORATORY  PRACTICE. 

may  be  readily  closed  by  melting  the  glass  or 
opened  by  breaking  off  the  tip  as  required.  Such 
a  bulb  tube  will  have  the  form  of  a  measuring  in- 
strument well  known  in  chemistry  as  a  pipette. 
The  tubulature  being  closed,  the  bulb  tube  is  first 
filled  with  colored  water  (freed  from  air  by  boiling 
in  an  open  beaker  glass)  to  a  point  about  30  milli- 
metres above  the  neck.  To  do  this,  first  heat  the 
empty  bulb  in  a  free  flame  (but  not  hotter  than  the 
hand  can  bear),  and  then  dip  the  open  end  of  the 
stem  into  the  colored  water.  Wait  until  a  teaspoon- 
f  ul  has  been  drawn  into  the  bulb,  and  then  bringing 
the  stem  upright  boil  carefully  the  water  in  the 
bulb  until  the  interior  is  full  of  steam  ;  then, 
grasping  the  bulb  with  some  holder  to  protect  the 
hand,  again  plunge  the  mouth  of  the  stem  in  the 
colored  water,  kept  boiling  meanwhile.  The  ex- 
periment requires  a  little  dexterity,  and  is  for  that 
very  reason  good  practice.  If  the  manipulation  is 
successful  the  bulb  tube  will  completely  fill  with 
water  without  showing  the  smallest  air  bubble. 
It  can  be  set  away  upright  when  still  nearly  boil- 
ing hot,  and  when  cold  the  water  will  be  found  to 
have  descended  in  the  stem  to  about  the  required 
amount,  provided  always  the  dimensions  given 
have  been  followed  in  the  construction  of  the 
bulbs. 

The  filled  bulb  is  now  to  be  packed  in  broken 
ice  and  the  level  to  which  the  liquid  sinks  in  the 


EXPANSION  OF  WATER  BY  HEAT.  21 

stem  marked  (most  readily  with  the  burnt  end  of 
a  match).  It  is  next  to  be  heated  in  a  steam  bath 
and  the  level  again  marked.  (The  bulb  must  not 
rest  on  the  overheated  bottom  of  the  steam  bath, 
else  large  fluctuations  of  level  will  be  noticed, 
caused  by  the  formation  of  steam  in  the  inte- 
rior.) 

In  order  now  to  measure  the  amount  of  expan- 
sion between  the  freezing  and  the  boiling  point 
the  bulb  tube  should  first  be  emptied,  breaking 
off  the  tip  of  the  tubulature  for  the  purpose. 
Using  now  the  instrument  as  a  pipette,  water 
should  be  sucked  into  the  bulb  and  tube  until  the 
curved  surface  (meniscus,  so  called)  is  tan- 
gent to  the  first  mark.  This  water  should  then  be 
run  out  into  a  small  tared  beaker  and  weighed. 
Again  the  bulb  and  tube  should  be  filled  but  to 
the  upper  mark,  and  then  the  column  of  water 
between  the  two  marks  run  out  with  the  greatest 
possible  care  into  a  small  tared  stoppered  bottle 
and  weighed  if  possible  to  a  milligramme.  It  is 
best  to  use  pure  distilled  water  for  the  purpose 
at  the  temperature  of  the  laboratory,  and  it  will 
require  a  little  experience  to  run  out  the  exact 
amount  of  water  between  the  marks.  But  re- 
peated trials  can  be  made  until  the  result  is  ob- 
tained, and  it  will  be  found  that  by  moistening 
the  finger  which  controls  the  mouth  of  the  stem 
great  sharpness  can  be  secured. 


22  LABORATORY  PRACTICE. 

As  the  two  weights  thus  obtained  are  to  each 
other,  so  is  the  volume  of  water  at  the  freezing 
point  to  the  increase  of  volume  at  the  boiling 
point.  Hence  it  will  be  easy  to  find  the  fraction 
of  its  volume  which  one  cubic  centimetre  of 
water  would  expand  between  these  temperatures, 
and  this  is  technically  called  the  coefficient  of 
expansion.  This  value  can  be  only  approxi- 
mately found  in  this  way ;  but,  as  in  the  first 
experiment,  so  here  and  in  all  subsequent  ex- 
periments involving  quantitative  measurement  the 
probable  error  with  the  tools  used  should  be  es- 
timated, although  the  point  may  not  again  be  re- 
ferred to. 

This  experiment  illustrates  not  only  the  expan- 
sion of  water  by  heat  and  its  contraction  when 
cooled,  but  also  the  expansion  of  air  by  heat,  as 
shown  in  the  method  of  filling  the  bulb.  The  fact 
that  it  is  the  relative  and  not  the  absolute  expan- 
sion of  water  which  is  measured  should  be  fully 
explained.  The  bulb  tube  should  be  kept  for 
another  experiment. 

(1)  The  coefficient  of  cubic  expansion  of  glass  (that  is, 
the  expansion  of  one  cubic  centimetre  in  volume  between 
0°  and  100°  C.)  is  0'0025.     What  is  the  absolute  expansion 
of  water  according  to  your  experiment  ? 

(2)  Why  is  it  important  to  determine  the  smaller  of  the 
two  weights  required  so  much  more  sharply  than  the  larger 
of  the  two  ? 

(3)  Why  boil  so  thoroughly  the  water  with  which  the 
bulb  tube  is  first  filled  ? 


MELTING  AND  BOILING  POINTS.  23 

Ex.  3.  (a)  The  Melting  and  Boiling  Points  of 
Water. — For  this  experiment  thermometers  with  a 
scale  on  the  tube  from  —5°  to  360°  (such  as  are  sold 
in  Germany  for  eighty  cents)  are  required.  The 
student  should  first  test  the  melting  point  of  ice, 
repeating  the  observation  several  times  with  differ- 
ent amounts  of  ice  and  under  different  conditions, 
until  he  gains  a  clear  idea  of  the  constancy  of  the 
thermal  state  at  which  the  change  takes  place. 
He  should  also  repeat  the  observation  with  dif- 
ferent thermometers,  using  if  possible  thermome- 
ters of  short  range,  with  the  degrees  divided 
into  tenths  ;  and  the  cause  of  what  is  called 
the  rise  of  the  zero  point  should  be  explained. 
Next,  the  boiling  point  of  water  should  be  ob- 
served, and  attention  should  be  called  to  the 
constancy  of  the  thermal  condition  at  which  this 
phenomenon  takes  place  under  the  same  atmos- 
pheric pressure  ;  and  the  small  variations  which 
depend  on  changes  of  the  barometer  should  also 
be  explained,  so  far  as  the  previous  knowledge 
of  the  student  will  permit.  The  chief  principle 
to  be  taught  by  this  experiment  is  the  constancy 
of  the  thermal  conditions  which  we  call  the  melt- 
ing point  of  ice  and  the  boiling  point  of  water. 

(1)  Does  the  quantity  of  ice  melting  or  the  quantity  of 
water  boiling-  make  any  difference  in  the  temperature  of 
these  materials  as  indicated  by  the  thermometer  ?  What 
inference  do  you  draw  from  these  facts  ? 


24  LABORATORY  PRACTICE. 

(2)  How  does  the  temperature  of  the  ice  compare  with 
that  of  the  lambent  water,  or  the  temperature  of  boiling 
water  with  that  of  the  steam  above  it  ?    What  becomes  of 
the  heat  that  enters  the  containing  vessel  ? 

(3)  On  what  evidence  do  you  infer  that  the  freezing  and 
boiling  points  of  water  are  constant  under  the  same  con- 
ditions ? 

(4)  What  is  the  effect  of  a  change  of  atmospheric  pressure 
on  the  boiling  point  ? 

(5)  It  is  often  the  case  when  a  thermometer  is  packed 
in  broken  ice  that  the  mercury  column  does  not  stand  ex- 
actly at  the  zero  point  of  the  scale.     What  is  the  cause  of 
this  anomaly,  and  ought  observations  made  with  the  instru- 
ment to  be  corrected  therefor  ?    When  the  thermometer  is 
dipped  in  water  to  which  an  inadequate  amount  of  ice  is 
added  the  thermometer  seldom  marks  the  true  zero.     Why 
is  this  to  be  expected  ? 

(b)  Distillation  of  Water. — Use  for  the  pur- 
pose a  glass  retort  holding  about  two  hundred 
cubic  centimetres.  First  try  to  distil  without 
any  provision  for  cooling  the  neck  of  the  retort. 
Afterwards  wrap  the  neck  loosely  with  several 
folds  of  filtering  paper,  so  as  to  form  a  covering 
(several  inches  shorter  than  the  neck)  secured  at 
the  two  ends  with  rubber  bands.  Adjust  a  hose 
connected  with  a  tap  so  that  water  will  trickle 
through  a  hole  made  for  the  purpose  near  the 
upper  edge  of  the  cover,  and  lead  off  the  stream 
from  the  lower  end  with  a  line  of  lamp  wicking 
that  has  been  previously  soaked  in  water.  Kepeat 
now  the  experiment. 

(1)  What  is  the  material  above  the  water  while  boiling 
in  the  retort  ?  Is  this  material  the  same  as  water  ?  Has  it 


PRINCIPLES  OF  THE  THERMOMETER.  25 

weight  ?  Is  there  necessarily  any  loss  of  weight  in  the  pro- 
cess of  distillation  ?  When  water  hoils  away  in  an  open 
kettle  what  "becomes  of  it  ? 

(2)  Why  is  it  necessary  to  adopt  some  means  of  cooling 
thf  retort  neck  in  order  to  distil  rapidly  ? 

Ex.  4.  Construction  and  Principles  of  the 
Thermometer. — The  student  may  next  make,  with 
the  tube  apparatus  of  Ex.  2,  a  water  thermome- 
ter, and  compare  it  with  one  of  the  ordinary  mer- 
cury thermometers  used  in  the  laboratory.  For 
this  purpose  the  tip  of  the  tubulature  must  be  re- 
sealed  and  the  bulb  filled  with  colored  water  with 
the  same  precautions  as  before.  The  fixed  points 
having  been  permanently  marked  (best  with  a  file 
wet  with  kerosene),  the  interval  between  them 
should  now  be  divided  into  twenty  equal  parts, 
corresponding  on  the  mercury  thermometer  to  5°, 
10°,  15°,  20°,  etc.,  and  a  pasteboard  scale  adjusted 
to  the  stem.  The  student  should  next  compare 
the  two  thermometers,  dipping  them  for  the  pur- 
pose side  by  side  and  gradually  raising  the  tem- 
perature of  the  bath.*  Let  him  note  the  heights 
of  the  water  column  corresponding  to  each  one  of 
the  series  of  degrees  of  the  mercury  thermometer 
just  mentioned,  and  mark  each  of  these  positions 

*  On  account  of  the  much  larger  mass  of  the  materials  to  be 
heated  and  the  much  greater  capacity  of  water  for  heat,  the  water 
thermometer  always  lags  behind  the  mercury  thermometer,  and  in 
order  that  the  comparison  should  be  just  the  temperature  must  be 
held  at  each  point  long  enough  for  an  equilibrium  to  be  established. 


26  LABORATORY  PRACTICE. 

on  the  scale.  After  the  positions  have  been 
marked  with  a  pencil,  the  scale  should  be  neatly 
drawn,  with  the  equal  divisions  on  one  side  and 
the  irregular  divisions  on  the  other  side  of  the 
stem. 

(1)  The  significance  of  this  experiment  will  not  be  ap- 
preciated unless  the  student  has  himself  made  a  mercury 
thermometer,  or  unless  its  construction  has  been  fully  ex- 
plained by  the  teacher.   There  are  difficulties  in  the  construc- 
tion which  do  not  recommend  it  as  a  general  laboratory  ex- 
periment, but  it  may  safely  be  entrusted  as  a  supplementary 
experiment  to  the  more  skilful  men.     Bulb  tubes  should  be 
provided  for  this  purpose  with  a  cup  at  the  mouth  of  the 
stem  to  hold  the  mercury  required  to  fill  the  bulb,  and  inex- 
perienced manipulators  will  be  most  successful  with  very 
small  bulbs  (not  holding  over  one  fifth  of  a  cubic  centi- 
metre) with  proportionally  fine  tubes. 

(2)  If  equal  intervals  on  the  scale  of  the  mercury  ther- 
mometer represent  equal  changes  of  temperature,  can  the 
same  be  true  of  the  water  thermometer  ? 

(3)  Using  only  the  second  scale  of  the  water  thermome- 
ter with  irregular  intervals  constructed  as  described  above 
(that  is,  adjusted  to  each  of  the  divisions  5°,  10°,  15°,  20°,  etc., 
of  a  mercury  thermometer),  it  is  obvious  that  the  water 
thermometer  would  give  the  same  indications  as  the  mer- 
cury thermometer ;  but  how  would  the  lengths  of  the  suc- 
cessive divisions  of  this  scale  vary  among  themselves  be- 
tween 0°  and  100°  ? 

(4)  What  must  be  the  cause  of  the  great  discrepancy 
between  the  two  thermometers,  when  both  are  graduated 
with  regular  intervals  ?  and  why  is  the  mercury  thermome- 
ter to  be  preferred  as  a  measure  of  temperature  ? 

(5)  Do  the  intervals  of  the  scale  of  the  mercury  ther- 
mometer correspond  to  equal  differences  of  temperature  ? 
And  what  must  be  the  nature  of  the  thermometric  material 
of  which  this  would  be  true  ?    Could  it  be  true  of  any  ma- 
terial in  a  glass  envelope  ?    How  is  it  with  air  ? 


POINT  OF   MAXIMUM  DENSITY.  27 

(6)  State  your  conception  of  the  term  temperature  as 
formed  from  this  experiment,  and  from  the  considerations 
suggested  by  the  above  questions. 

(7)  It  is  not  expected  that  the  student  can  fully  answer 
all  these  questions  from  his  previous  knowledge.     They  are 
intended  to  stimulate  thought  and  suggest  to  the  teacher 
directions  in  which  the  subject  may  be  developed.     The  stu- 
dent can  at  least  be  made  to  comprehend  that  temperature 
is  a  thermal  condition  of  which  equal  intervals  can  be  predi- 
cated and  measured  with  close  approximation,  although  the 
condition  itself  can  only  be  denned  theoretically. 

Ex.  5.  (a)  Irregular  Expansion  of  Water 
near  its  Freezing  Point.  —  Having  packed  in 
broken  ice  the  water  thermometer  described  under 
the  last  experiment,  wait  until  the  column  is  sta- 
tionary, and  then  raise  the  instrument  from  the 
ice  and  watch  the  motion  of  the  column  as  the 
temperature  rises.  Note  that  the  column  sinks  to 
a  perceptible  extent  before  it  begins  to  rise,  but 
afterwards  expands  steadily  with  the  increasing 
temperature.  The  experiment  is  intended  to  show 
that  the  point  of  the  maximum  density  of  water 
is  above  the  freezing  point,  and  in  connection  with 
it  the  relations  of  this  remarkable  property  of 
water  in  the  economy  of  nature  should  be  ex- 
plained to  the  student. 

(b)  Conduction  of  Water  for  Heat. — Select  a 
narrow  but  long  test  tube,  nearly  fill  with  water* 
Grasping  the  tube  at  the  bottom,  apply  a  flame 
near  the  top  of  the  tube  a  few  centimetres  from 
the  surface  of  the  water  until  the  liquid  boils. 


28  LABORATORY  PRACTICE. 

Repeat  the  experiment,  applying  the  flame  near 
the  bottom  of  the  tube  while  holding  it  at  the 
top. 

(c)  Conduction  Currents.  —  Repeat  the  last 
phase  of  the  experiment,  adding  some  shreds  of 
filtering  paper  to  the  water,  whose  motion  will  in- 
dicate the  play  of  the  currents. 

(1)  Why  can  water  be  most  readily  heated  by  applying 
the  flame  at  the  bottom  of  the  containing  vessel  ? 

(2)  Why  in  winter  do  the  ponds  only  freeze  on  the  sur- 
face ? 

(3)  Would  a  water  thermometer  made  as  above  show 
the  exact  point  of  maximum  density  ? 

Ex.  6.  (a)  Density  of  Ice.  First  Method.— 
Float  a  lump  of  ice  of  regular  shape  on  water  and 
observe  as  nearly  as  you  can  estimate  with  the  eye 
what  fraction  of  the  volume  is  immersed.  Mix 
now  about  an  equal  volume  of  alcohol  with  the 
water  until  the  ice  neither  floats  nor  sinks,  and 
then,  removing  the  ice,  take  the  specific  gravity 
of  the  mixture  by  means  of  the  tin  cylinder  of 
Ex.  1. 

(b)  Second  Method. — Provide  a  tin  cylindrical 
vessel  of  the  capacity  of  about  250  cubic  centime- 
tres, which  should  be  packed  round  with  ice  and 
salt  like  an  ice-cream  freezer  (use  glass  beaker  500 
cubic  centimetres  for  outer  vessel) ;  provide  a  sec- 
ond, similar  in  every  respect,  but  two  thirds  filled 
with  kerosene.  Suspend  now  a  piece  of  ice  weigh- 


DENSITY  OF  ICE,  29 

ing  about  twenty-five  grammes  to  the  beam  of  the 
balance  by  means  of  a  silk  thread,  as  in  Ex.  1, 
and  adjust  so  that  the  ice  shall  hang  within  the 
first  cylinder,  and  weigh  the  ice.  Replacing  then 
the  first  cylinder  with  the  second,  weigh  the  ice 
immersed  in  kerosene.  Lastly,  weigh,  immersed 
in  kerosene,  the  metallic  cylinder  described  in 
Ex.  1.  We  have  now  four  weights — the  weight 
of  ice  in  air,  the  weight  of  the  same  ice  immersed 
in  kerosene,  the  weight  of  the  metallic  cylinder 
immersed  in  the  same  kerosene,  and  from  Ex.  1 
we  have  the  weight  of  this  cylinder  immersed  in 
ice-cold  water.  From  these  data  we  can  easily 
calculate  the  specific  gravity  and  density  of  the 
ice. 

(1)  The  second  method  requires  more  care  as  well  as 
more  apparatus  than  the  first,  and  may  be  reserved  for  the 
more  skilful  experimenters. 

(2)  What  fraction  of  an  iceberg  is  immersed  in  the 
ocean  ?    Specific  gravity  of  sea  water,  1*03. 

(3)  A  block  of  ice  weighs  36  '72  kilogrammes.     What  is 
its  volume  ? 

Ex.  7.  Expansion  of  Water  in  Freezing. 
—Provide  glass  bulb  tubes  similar  in  construc- 
tion to  those  described  under  Ex.  2,  only  hav- 
ing larger  stems,  about  three  millimetres  in  diame- 
ter. Fill  the  bulb  about  one  half  with  water  and 
boil  the  liquid  to  expel  the  air.  Fill  the  rest  of 
the  bulb  and  a  small  portion  of  the  stem  with 
kerosene.  (This  is  easily  .done  with  a  small  tun- 


30  LABORATORY  PRACTICE. 

nel,  the  neck  of  which  has  been  drawn  out  into  a 
long  fine  tube.)  Next  freeze  the  water  by  im- 
mersing the  bulb  in  a  mixture  of  ice  and  salt.  To 
prevent  breaking  the  glass,  hold  the  stem  oblique- 
ly and  keep  the  bulb  turning  while  the  water  is 
freezing.  Raise  the  tube  from  the  freezing  mixt- 
ure and  follow  the  descent  of  the  column  as  the 
ice  melts.  The  motion  of  the  column  is  somewhat 
erratic,  owing  to  the  circumstances  that  the  water, 
kerosene,  and  glass  expand  independently  and  at 
very  different  rates,  and  that  the  heat  diffuses 
through  those  materials  only  slowly ;  but  the  main 
feature  of  the  experiment  far  surpasses  all  the 
lesser  effects.  Observe  all  variations  as  closely  as 
possible  and  seek  to  explain  them. 

(1)  The  volume  of  water  used  in  the  bulb  can  readily  be 
found  by  weighing  the  glass  before  adding  the  water  and 
after  the  water  has  been  boiled  and  cooled.     The  expansion 
of  the  water  in  freezing  can  then  be  measured  by  following 
the  same  general  method  described  under  Ex.  2  and  the  re- 
sult compared  with  those  of  the  previous  experiment.     This 
phase  of  the  experiment  may  serve  when  occasion  offers  as 
an  extra  exercise. 

(2)  Why  boil  the  water  before  freezing  ? 

(3)  Any  attempt  to  measure  the  density  of  steam  would 
be  premature  in  this  connection,  but  the  general  result  of 
such  measurements  should  here  be  stated  and  illustrations 
given  of  the  great  expansion  which  the  conversion  of  water 
into  steam  involves.     The  specific  gravity  of  steam  is  0'6235 
referred  to  air  at  the  same  temperature  and  pressure,  and 
one  cubic  centimetre  of  boiling  water  yields  approximately 
1,627  cubic  centimetres  of  free  steam  at  the  normal  pressure 
of  the  air  (one  cubic  inch  yields  nearly  one  cubic  foot).     In  a 


CAPACITY  FOR  HEAT.  31 

closed  vessel  partially  filled  with  water— like  a  steam  boiler— 
both  the  density  and  pressure  of  the  steam  which  forms  in 
the  assumed  empty  space  above  the  liquid  vary  with  the 
temperature,  increasing  rapidly  as  the  temperature  rises,  ac- 
cording to  a  complex  law.  Tables  giving  the  density  and 
pressure  corresponding  to  successive  temperatures  are  to  be 
found  in  works  on  physics,  and  as  these  values  have  impor- 
tant applications  even  in  this  elementary  book  they  should  be 
shown  and  explained.  When  high-pressure  steam  escapes 
from  a  boiler  into  the  atmosphere,  the  steam  expands  until 
its  pressure  is  reduced  to  that  of  the  atmosphere,  and  then 
mixes  with  the  air.  A  locomotive  engine  is  constantly  pump- 
ing steam  into  the  atmosphere,  and  at  every  revolution  of  the 
driving  wheels  each  cylinder  is  filled  and  emptied  twice. 
In  a  run  of  twenty  miles  a  very  large  volume  of  steam  is 
thus  spent,  and  is  supplied  by  the  evaporation  in  the  boiler 
of  a  comparatively  small  amount  of  water. 

Ex.  8.  Capacity  of  Water  for  Heat. — Provide 
a  round  pasteboard  box  ;  line  the  inside  with  felt 
at  least  two  inches  thick,  covering  both  the  bot- 
tom of  the  box  and  the  under  side  of  the  cover, 
but  not  the  rim,  which  should  be  at  least  three 
inches  wide ;  cover  the  felt  with  a  second  lining 
of  thick  paper.  Procure  also  a  cylindrical  vessel 
made  of  the  thinnest  sheet  brass  which  can  be 
obtained.  This  vessel  should  be  ten  centimetres 
in  diameter  and  fifteen  centimetres  high,  with 
thin  flanges  on  the  sides  and  bottom.  It  is  to 
stand  inside  the  felt-lined  box,  from  which  it  is 
kept  apart  by  the  flanges,  and  the  dimensions 
should  be  such  as  to  leave  half  an  inch  of  air 
space  around  the  brass  vessel.  It  should  be  pro- 
vided with  a  stirrer,  made  also  of  thin  sheet  brass, 


32  LABORATORY   PRACTICE. 

like  a  turbine  wheel ;  and  there  must  be  a  hole 
through  the  felt-lined  cover  of  the  outer  box  large 
enough  to  pass  the  tubular  axis  of  the  stirring 
wheel,  and  the  thin  brass  tube  which  forms  the 
axis  must,  in  its  turn,  be  large  enough  to  pass 
the  bulb  of  a  thermometer.  This  apparatus  is 
called  a  calorimeter,  and  will  be  used  for  many 
experiments. 

Weigh  out  in  the  brass  vessel  about  500 
grammes  of  water,  noting  the  exact  weight.  Re- 
place in  the  calorimeter,  and  wait  until  the 
temperature  is  constant,  as  indicated  by  a  ther- 
mometer graduated  to  one  tenth  of  a  centigrade 
degree.  Weigh  out  next  in  a  dipper,  or  some 
other  vessel  with  a  handle  which  can  convenient- 
ly be  heated  in  a  steam  bath  (a  farina  kettle, 
for  example),  about  500  grammes  of  small  iron 
nails.  When  the  metal  has  fully  reached  the 
temperature  of  the  steam  (and  this  will  be  most 
readily  secured  by  covering  the  top  of  the  kettle 
with  a  towel)  pour  the  nails  as  quickly  as  pos- 
sible into  the  water,  which  may  be  left  uncovered 
during  this  experiment.  Follow  now  the  tem- 
perature of  the  water,  and  note  the  highest  point 
reached,  taking  care  to  secure  uniformity  by  mov- 
ing the  stirrer  several  times  up  and  down  before 
each  reading  of  the  thermometer.  The  tempera- 
ture will  rise  only  a  few  degrees  ;  for  the  quantity 
of  heat  given  out  by  the  metal  in  cooling  from 


CAPACITY  FOR  HEAT.  33 

100°  to  the  temperature  of  the  calorimeter  is  only 
sufficient  to  raise  the  temperature  of  the  water  by 
a  comparatively  small  amount,  showing  that  the 
capacity  of  the  water  for  heat  is  far  greater  than 
that  of  the  iron  nails.  Make  now  a  similar  ex- 
periment with  granulated  copper  or  with  brass 
turnings,  a  third  with  lead  shot,  and  a  fourth 
with  glass  beads.  By  means  of  the  data  thus 
obtained  the  student  should  calculate  the  ca- 
pacity of  iron,  copper,  lead,  and  glass  for  heat, 
as  compared  with  water.  Adopting  for  the  unit 
of  measure  that  quantity  of  heat  which  will 
raise  the  temperature  of  one  gramme  of  water 
one  centigrade  degree,  the  student  will  find  what 
fraction  of  a  unit  of  heat  will  raise  the  tempera- 
ture of  a  gramme  of  copper,  of  iron,  of  lead,  or  of 
glass  one  degree.  These  values  are  called  the 
specific  heats  of  the  substances.  It  is  expected 
that  the  student  will  gain  through  these  measure- 
ments a  clear  conception  of  the  difference  between 
temperature  and  quantity  of  heat;  and  this  is 
the  most  important  point  here  illustrated.  After 
the  student  fully  understands  how  temperature 
is  measured  and  how  quantity  of  heat  is  meas- 
ured, he  will  be  easily  able  to  construct  the  sim- 
ple formulae  by  which  the  ordinary  problems  con- 
nected with  specific  heat  are  solved.  These  prob- 
lems should  be  multiplied  until  the  subject  is 
fully  mastered.  The  student  will  thus  come  to 


34  LABORATORY  PRACTICE. 

appreciate  how  very  great  is  the  capacity  of  water 
for  heat,  and  he  should  be  shown  the  important 
relations  of  this  storage  capacity  in  the  economy 
of  nature. 

(1)  It  is  an  obvious  caution  in  this  experiment  to  pre- 
vent the  calorimeter  from  being  affected  by  radiation  from 
the  steam  bath. 

(2)  Make  clear  the  distinction  between  the  unit  of  heat 
and  the  unit  of  temperature, 

(3)  How  many  units  of  heat  would  be  required  to  raise 
the  temperature  of  100  grammes  of  brass  from  18°  C.  to 
23°  C.  ?    How  many  grammes  of  water  would  be  the  ther- 
mal equivalent  in  this  respect  of  100  grammes  of  brass  ? 

(4)  Must  not  the  brass  vessel  of  the  calorimeter  have  an 
influence  on  the  results  of  the  above  experiments  ?    How 
can  you  find  its  thermal  equivalent  and  correct  your  results 
therefor  ?    Note  that  this  correction  must  always  be  made 
in  all  similar  cases,  as  in  the  two  following  experiments. 

(5)  Why  are  insular  climates  comparatively  moderate  ? 

Ex.  9.  Latent  Heat  of  Water.—  Prepare  the 
calorimeter  as  in  Ex.  8,  and  note  the  weight  and 
temperature  of  the  water.  Stir  in  finely  broken 
ice  so  long  as  it  promptly  melts.  Close  now  the 
the  calorimeter  and  note  the  fall  of  temperature. 
Lastly,  remove  and  reweigh  the  brass  vessel ;  and 
this  weight,  when  compared  with  the  first  weight, 
will  give  the  amount  of  ice  added.  If  no  heat 
were  required  to  melt  the  ice,  it  is  obvious  that 
its  effect  on  the  calorimeter  would  have  been  the 
same  as  that  of  an  equal  quantity  of  ice-cold 
water.  Bearing  this  in  mind,  it  will  be  easy  to 
calculate  from  the  experimental  data  how  many 


LATENT  HEAT.  35 

units  of  heat  are  required  to  melt  one  gramme  of 
ice.  This  is  called  the  latent  heat  of  water.  Ex- 
plain the  nature  of  this  phenomenon  and  the  ob- 
jections to  the  use  of  the  term  "latent"  as  applied 
to  it. 

(1)  Take  care  that  the  broken  ice  is  as  dry  as  possible 
and  not  drenched  with  running  water.     Why  ? 

(2)  How  many  grammes  of  broken  ice  would  be  required, 
when  stirred  into  500  grammes  of  water  at  20°  C.,  to  reduce 
the  temperature  of  the  liquid  to  0°  C.  ?    Why  should  you 
expect  that  practically  more  than  the  theoretical  amount 
would  be  required  ? 

(3)  In  what  sense  can  freezing  be  regarded  as  a  warming 
process  ? 

Ex.  10.  Latent  Heat  of  Steam. — Prepare  the 
calorimeter  as  before,  and  pass  into  the  water  for 
a  few  minutes  a  current  of  dry  steam.  The  steam 
may  be  drawn  from  the  steam  pipes  of  the  labora- 
tory through  tubes  of  thin  sheet  brass,  which,  if 
first  warmed  by  passing  the  current  for  a  few 
moments  before  dipping  the  mouth  of  the  tube 
under  the  water,  will  deliver  nearly  dry  steam. 
When  the  laboratory  is  not  heated  by  steam  the 
steam  may  be  generated  from  a  glass  flask ;  but 
this  must  be  so  screened  as  not  to  affect  the  calo- 
rimeter. Stir,  and  observe  the  rise  in  tempera- 
ture. By  reweighing  the  brass  vessel  the  amount 
of  steam  condensed  is  determined ;  and  by  apply- 
ing the  same  course  of  reasoning  as  in  the  last 
experiment  the  student  can  easily  calculate  the 


36  LABORATORY  PRACTICE. 

amount  of  heat  developed  when  one  gramme  of 
steam  condenses  to  one  gramme  of  boiling-hot 
water  without  change  of  temperature.  We  thus 
measure  the  latent  heat  of  free  steam — that  is,  of 
such  steam  as  rises  from  water  boiling  under  the 
ordinary  pressure  of  the  atmosphere,  which  neces- 
sarily has  the  same  tension  as  the  atmosphere 
and  the  same  temperature  as  the  boiling  water. 
The  teacher  may  here  add  that  the  latent  heat  of 
water  changes  with  the  temperature  according  to 
a  well-known  law.  Let  him  also  consider  in  this 
connection  the  use  of  steam  in  heating,  and  also 
the  effect  of  the  aqueous  circulation  in  modifying 
the  temperature  of  the  globe. 

(1)  Why  must  the  steam  be  dry  ? 

(2)  Why  the  necessity  of  cooling  the  neck  of  the  retort 
In  the  experiment  on  the  distillation  of  water  ? 

(3)  Why  do  the  rain  storms  tend  to  equalize  the  climates 
of  the  earth  ? 

(4)  In  many  industrial  processes  water   is   boiled   in 
wooden  tanks  by  blowing  steam  into  the  liquid.     In  such 
a  case,  if  we  start  with  1,000  kilogrammes  of  water  at  15°, 
how  much  liquid  (condensed  steam)  will  be  added  in  rais- 
ing the  temperature  to  the  boiling  point  ?     It  is  here  as- 
sumed that  there  is  no  loss,  but  as  the  water  nears  boiling 
a  great  deal  of  steam  escapes.     How  does  this  loss  affect  the 
result  ? 

Ex.  11.  Solvent  Power  of  Water.— \.  Weigh 
into  a  test  tube  ten  grammes  of  water.  Weigh 
out  successive  portions  of  one  gramme  each  of 
cupric  sulphate.  Add  the  first  portion  to  the 


SOLVENT  POWER  OF   WATER.  3f 

test  tube,  cork  tightly,  and  shake  ;  and  if  this  dis- 
solves, add  the  second,  and  so  on  until,  after  con- 
tinned  shaking,  the  last  portion  fails  to  dissolve. 
Estimate  the  number  of  parts  of  the  salt  that  have 
dissolved  in  one  hundred  parts  of  water.  In  like 
manner  test  the  solubility  of  potassic  nitrate,  po- 
tassic  sulphate,  acid  potassic  tartrate,  and  baric 
sulphate.*  If  even  the  first  portion  fails  to  dis- 
solve, ascertain  whether  any  has  dissolved.  This 
is  best  done  by  filtering  some  of  the  mixture  and 
evaporating  a  few  drops  of  the  filtrate  (which 
should  be  perfectly  clear)  on  a  strip  of  window 
glass. 

2.  In  the  same  way  experiment  with  common 
salt  (sodic  chloride),  baric  nitrate,  and  crystallized 
sodic  sulphate ;  but  when  each  solution  has  be- 
come saturated  at  the  temperature  of  the  room, 
warm  slowly,  and  as  often  as  the  salt  is  entirely 
dissolved  add  a  new  portion  of  one  gramme,  final- 
ly bringing  the  liquid  to  boiling.  Carefully  ob- 
serve the  effects.  What  does  the  experiment 
show?  What  effect  may  the  heat  of  the  hand 
produce  in  the  examples  under  1  ? 

Points  to  be  made  prominent :  The  very  gen- 
eral, but  limited,  solvent  power  of  water ;  effect 

*  In  cases  like  this,  where  a  repetition  of  the  same  experiment 
would  be  tedious,  and  the  additional  practice  not  important,  the 
end  will  be  gained  by  having  the  determinations  under  different 
conditions  made  by  different  students,  and  the  results  compared  be- 
fore the  class. 


38  LABORATORY  PRACTICE. 

of  temperature  on  the  solvent  power ;  features 
exhibited  by  different  substances,  both  as  regards 
the  extent  of  their  solubility  and  its  variation 
with  the  temperature ;  the  conception  of  a  satu- 
rated solution. 

(1)  When  hot  water  dissolves  more  of  a  salt  than  cold, 
what  should  you  expect  would  follow  on  cooling  a  hot  satu- 
rated solution  ?    Try  the  experiment  with  nitre  and  copper 
sulphate. 

(2)  If  water  holds  a  non-volatile  material  in  solution, 
what  should  you  anticipate  would  follow  on  evaporating 
the  water  ?    What  if  the  material  were  in  itself  volatile  ? 
Try  the  experiment  by  evaporating  on  a  glass  plate  a  few 
drops  of  solution  :  (1)  of  common  salts,  (2)   of  chloride  of 
ammonia,  and  (3)  of  aqua  ammonia,  using  in  each  case  weak 
solutions  and  heating  cautiously  on  iron  plate  over  a  flame. 

Ex.  12.  Hard  Water.— Boil  down  in  a  clean 
saucepan  half  a  litre  of  well  water  to  a  few  cubic 
centimetres.  Transfer  to  a  tared  porcelain  cruci- 
ble, and,  after  evaporating  to  dryness,  weigh  the 
residue  and  try  to  recognize  the  chief  ingredient 
by  the  taste.  Calculate  the  per  cent  of  solid  im- 
purities dissolved  in  the  well  water.  Redissolve 
the  residue  as  far  as  possible  in  a  few  cubic  centi- 
metres of  water,  and  shake  up  in  a  test  tube  with 
a  solution  of  soap,  adding  the  soap  solution  in 
small  successive  portions.  Distil  another  portion 
of  the  same  well  water,  and  test  the  distillate  by 
evaporating  a  few  drops  on  a  glass  plate,  and  with 
soap  solution  as  before.  Compare  the  results. 

Ex.  13.  The  Crystallizing  Power  of  Water.  — 


MEDIUM  OF  CHEMICAL  CHANGES.  39 

Prepare  a  saturated  solution  (about  20  grammes 
each)  of  alum,  potassic  ferrocyanide,  potassic  ni- 
trate, sodic  nitrate,  ferrous  sulphate,  and  cupric 
sulphate.  Pour  each  solution  into  a  shallow  dish, 
and,  protecting  it  from  dust  with  a  cover  of  po- 
rous paper,  leave  the  solution  to  evaporate  in  a 
warm,  dry  place.  Examine  from  time  to  time, 
and  study  the  forms  obtained.  This  work  may  be 
divided  to  advantage  among  several  students,  who 
may  be  shown  how  to  "  nurse  "  the  crystals,  and 
will  compete  with  each  other  to  obtain  large  and 
perfect  forms.  In  this  connection  the  fundamen- 
tal forms  of  crystals  should  be  studied,  and  the 
production  of  natural  crystals  in  geodes  and  min- 
eral veins  explained.  We  have  selected  one  ex- 
ample from  each  system  of  crystals  ;  but  the 
teacher  may  multiply  these  examples  according  to 
the  material  at  his  command,  so  as  to  illustrate  all 
the  characteristic  features  of  crystalline  growth. 

(1)  May  not  crystals  be  obtained  as  readily  by  cooling  a 
hot  saturated  solution  as  explained  above.     What  advan- 
tage is  to  be  gained  with  the  slower  process  recommended 
here  ? 

(2)  Describe  the  six  types  of  crystals  exhibited  by  the 
substances  above  selected. 

(3)  Are  these  crystalline  forms  accidental,  depending  on 
external  relations,  or  are  they  qualities,  and  therefore  char- 
acteristic of  the  substances  used  ? 

Ex.  14.  (a)   Water  the  Medium  of  CTiemical 
Changes. — Mix  in  a  mortar  half  a  gramme  of  pul- 


40  LABORATORY  PRACTICE. 

verized  dry  sodic  bicarbonate  with  the  same 
weight  of  pulverized  dry  tartaric  acid.  Transfer 
to  a  test  tube  and  pour  on  water.  After  the  ac- 
tion has  subsided  evaporate  the  liquid  to  dryness 
and  compare  by  tasting  the  residue  with  the  mate- 
rials used. 

(b)  Mix  in  a  mortar  a  few  milligrammes  of  dry 
pulverized  baric  chloride  with  about  the  same 
amount  of  dry  pulverized  sodic  sulphate  and  ob- 
serve the  effect.  Weigh  out  now  as  accurately  as 
possible  208  milligrammes  of  the  first  salt  and  142 
milligrammes  of  the'  second.  Dissolve  each  in 
separate  test  tubes  in  about  five  cubic  centimetres 
of  water.  Heat  the  solutions  nearly  to  boiling, 
and  pour  the  first  into  the  second.  When  cool 
enough  to  handle,  shake  the  materials  together 
and  leave  to  settle.  Pour  off  (decant)  a  portion 
of  the  clear  liquid  and  evaporate  it  to  dryness. 
Taste  the  residue  and  compare  with  the  substances 
taken. 

(1)  Why  weigh  so  accurately  the  materials  used  ? 

(2)  The  chief  point  to  be  noticed  in  the  above  experi- 
ments is  that  the  dry  powders  are  perfectly  inert  towards 
each  other  and  that  no  change  takes  place  until  water  is 
added.     In  saying  that  water  is  the  medium  of  the  change 
we  mean  that  it  acts  chiefly  in  virtue  of  its  solvent  power,  al- 
though it  often  happens  (as  will  afterwards  appear)  that  a 
portion  of  the  water  present  may  enter  into  union  with  the 
product  formed  or  may  be  separated  from  the  materials  as 
one  of  the  results  of  the  process.     This  concurrence  of  water 
in  chemical  changes  is  so  universal  as  to  be  one  of  the  most 


WATER  IN  COMBINATION.  41 

important  features  of  such  processes  ;  and  as  it  very  fre- 
quently obscures  more  essential  phases,  the  student  should 
become  acquainted  with  the  general  facts  from  the  start. 

(3)  The  solid  precipitate  which  falls  in  (6)  (baric  sul- 
phate) is  obviously  insoluble  in  water,  and  when  materials 
are  brought  together  in  solution  there  is  always  a  tendency 
to  such  a  transfer  of  their  constituent  parts  as  will  produce 
insoluble  compounds.  The  effect  in  this  case  is  one  of  wide 
significance  and  important  application. 

Ex.  15.  Water  in  Combination. — Heat  some 
small  bits  of  tlie  mineral  gypsum  at  the  bottom  of 
a  closed  glass  tube  held  obliquely,  and  seek  to 
recognize  the  clear  liquid  drops  which  condense 
on  the  walls.  Heat  now  a  weighed  amount  of 
gypsum  in  a  porcelain  crucible,  and  from  the 
weight  of  the  residue  determine  what  per  cent  of 
water  gypsum  contains.  Make  similar  experi- 
ments with  crystallized  baric  chloride.  Try  also, 
for  comparison,  an  experiment  with  common  salt, 
using  only  one  gramme  of  finely  pulverized  mate- 
rial and  covering  the  crucible  to  avoid  loss  by 
snapping.  As  illustrating  the  same  points,  take 
next  a  lump  of  quicklime  weighing  about  fifty 
grammes,  and,  having  noted  the  exact  weight, 
place  it  in  a  capacious  evaporating  dish  previously 
tared.  Now  pour  upon  it  water  little  by  little  so 
long  as  the  liquid  is  absorbed  and  weigh  the  prod- 
uct. From  these  weights  it  will  be  seen  that 
water  has  united  with  the  lime.  For  the  last 
illustration,  mix  with  water  to  a  thin  paste  some 
plaster  of  Paris  (dried  gypsum),  and  pour  the 


42  LABORATORY  PRACTICE. 

plaster  over  a  silver  dollar  previously  oiled  and 
placed  on  a  sheet  of  oiled  paper.  When  the  plas- 
ter has  hardened  remove  the  cast.  In  this  con- 
nection the  teacher  should  define  the  term  "  water 
of  crystallization"  and  make  the  distinction  be- 
tween analysis  and  synthesis. 

(1)  How  can  you  prove  that  the  liquid  drops  above  men- 
tioned are  water  ? 

(2)  Which  of  the  above  processes  are  analysis  and  which 
synthesis  ? 

Ex.  16.  Water  itself  a  Compound. — At  this 
point  the  student  should  be  shown  the  decomposi- 
tion of  water  by  an  electric  current  and  its  subse- 
quent synthesis  with  the  eudiometer  in  order  that 
he  may  clearly  see  that  our  knowledge  that  water 
is  composed  of  oxygen  and  hydrogen  gases  rests 
on  evidence  of  the  same  kind  as  that  which  ap- 
peared in  the  previous  experiment. 

These  experiments  can  not  be  performed  by 
the  student  himself  without  more  expensive  ap- 
pliances than  most  schools  can  command,  and  are 
best  shown  to  the  whole  class  at  once  on  the 
lecture  table.  The  writer  would  here  state  that  in 
his  opinion  it  is  not  necessary,  in  order  to  secure 
the  full  advantages  of  the  experimental  method, 
that  each  student  should  perform  every  experi- 
ment for  himself.  Indeed,  if  this  is  attempted,  a 
course  in  chemistry  must  either  be  made  very 
meagre  or  very  expensive.  If  the  student  actu- 


AIR  HAS  WEIGHT.  43 

ally  performs  in  the  laboratory  a  sufficient  num- 
ber of  experiments  to  give  him  the  spirit  of  the 
method,  he  will  usually  comprehend  the  full  sig- 
nificance of  those  which  are  plainly  exhibited  on 
the  lecture  table  by  the  instructor. 

2.  Air  as  an  Example  of  Aeriform  Matter. 

Ex.  17.  Air  has  Weight.— Select  two  flasks  of 
about  250  cubic  centimetres  capacity  which  have 
been  blown  in  the  same  mould  and  are  therefore 
of  equal  size.  Fit  them  tightly  with  corks.  Cork 
one  permanently  and  reserve  it  as  a  counterpoise. 
Add  to  the  other  about  twenty -five  cubic  centi- 
metres of  water  and  boil  the  water  over  a  lamp 
until  the  interior  is  full  of  steam ;  then  cork,  re- 
move the  lamp,  and  allow  to  cool.  Tare  now  the 
second  flask  with  the  first  and  such  additional 
weight  as  may  be  necessary.  Draw  the  cork  from 
the  second  flask,  and  after  air  has  entered  deter- 
mine the  increase  of  weight.  Determine  the  vol- 
ume of  the  air  weighed  by  adding  water  from  a 
graduated  measure  until  the  flask  is  filled  to  the 
former  level  of  the  cork.  From  these  values  the 
weight  of  one  litre  of  air  under  the  conditions  of 
the  experiment  can  be  roughly  determined. 

(1)  Accurate  methods  of  determining  the  weight  of  aeri- 
form substances  are  beyond  the  reach  of  elementary  students, 
but  the  values  obtained  for  a  few  of  the  gases  referred  to  in 
this  book  may  here  be  collected  for  reference. 


44  LABORATORY  PRACTICE. 

Weight  of  one  litre  when  H  =  760  millimetres  and  T  = 
0°.  Also  specific  gravity  when  air  =  1 : 

Weight.  Specific  Gravity. 

Air 1-2932  1* 

Oxygen 1*4303  1*1056 

Nitrogen 1*2562  0*9713 

Hydrogen 0*0896  0*0692 

Carbonic  dioxide 1*9775  1*5291 

Nitrous  oxide 1*9746  1*5269 

(2)  Why  is  it  so  much  more  difficult  to  determine  the 
weight  of  a  mass  of  gas  than  that  of  a  solid  or  liquid  body  ? 
Is  it  a  definite  quantity  ? 

(3)  What  is  the  use  of  the  corked  flask  as  a  part  of  the 
counterpoise  ? 

Ex.  18.  Relation  of  Volume  to  Pressure. 
(Law  of  Mariotte.) — Required  a  graduated  glass 
tube  about  300  millimetres  long,  divided  into  fifths 
of  a  cubic  centimetre,  open  at  the  lower  end  and 
closed  at  the  upper  by  a  tubulature  guarded  by  a 
pinch  cock  like  a  Mohrs  burette,  also  a  piece  of 
plain  glass  tubing  of  about  the  same  dimensions 
as  the  first.  Connect  the  two  with  a  length  (about 
300  millimetres)  of  stout  rubber  tubing  and  wire 
the  ends  on  to  the  open  mouths  of  the  glass  tubes 
and  then  support  with  clamps  the  two  tubes  side 
by  side,  the  rubber  connecting  tube  hanging  in  a 
curve  below  (the  last  must  be  stout  enough  to  pre- 
vent kinking).  Open  now  the  tubulature  and 
pour  mercury  into  the  open  mouth  of  the  side 
tube  until  the  level  of  the  two  columns  (which 
will  be  the  same  if  there  is  no  obstruction)  stands 
about  midway  of  the  graduation.  Close  the  tubu- 


LAW  OF  MARIOTTE.  45 

lature  and  read  the  volume  of  the  confined  air ; 
also  read  the  barometer,  which  gives  in  millime- 
tres of  mercury  column  the  pressure  to  which  the 
confined  air  is  exposed.  Next  raise  or  lower  the 
side  tube  to  as  great  an  extent  as  the  apparatus 
will  allow,  but  only  a  small  amount  at  a  time,  and 
in  each  position  measure  with  a  graduated  rule 
the  difference  in  the  heights  of  the  two  mercury 
columns  in  millimetres.  Read  the  corresponding 
volume  of  the  air,  and  also  take  again  the  height 
of  barometer  if  there  has  been  any  change.  Add 
to  each  difference  of  level  (positive  or  negative) 
the  corresponding  height  of  the  barometer.  Call 
this  sum  H,  which  stands  for  a  number  of  milli- 
metres. Make  out  a  table  giving  on  one  side  the 
observed  volumes,  Y,  in  cubic  centimetres,  and  on 
the  other  side  in  a  parallel  column  the  values  of  H 
in  millimetres.  Look  now  for  a  relation  between 
the  quantities  thus  tabulated,  and  it  should  be 
found  that  in  each  case 

V  :  V  =  H' :  H. 

(1)  A  Mohrs  burette  can  be  used  for  this  experiment, 
only  as  it  is  graduated  the  wrong  way  the  value  of  the  space 
between  the  lowest  division  and  the  stop  cock  must  first  be 
measured  and  added  to  the  readings  reversed. 

(2)  In  handling  the  apparatus  the  student  must  be  very 
careful  not  to  heat  the  measuring  tube,  but  give  time  before 
reading  for  the  whole  to  come  to  the  temperature  of  the 
room,  which  is  assumed  to  be  constant  during  the  experi- 
ment. 

(3)  This  experiment  is  not  only  calculated  to  give  a  clear 

4 


46  LABORATORY   PRACTICE. 

conception  of  Mariotte's  law,  but  it  also  furnishes  important 
practice  in  the  measurement  of  gas  volumes  and  gas  ten- 
sion. It  assumes  a  knowledge  of  the  use  of  the  barome- 
ter and  of  the  fundamental  principles  of  pneumatics,  and, 
like  these  subjects,  might  be  relegated  to  the  course  on 
physics.  But  the  whole  subject  is  of  such  fundamental 
importance  in  the  theory  of  chemistry  that  the  chemical 
student  can  not  fail  to  profit  by  the  exercise,  even  if  it  is 
a  review  of  previous  work. 

(4)  What  is  meant  by  the  tension  of  a  gas,  and  how  is  it 
measured  ?    Is  there  any  distinction  between  tension  and 
pressure  ?    What  is  the  standard  pressure  to  which  all  meas- 
urements of  gas  volumes  should  be  reduced  ? 

(5)  A  volume  of  air  was  found  to  be  200  cubic  centi- 
metres.    The  barometer  at  the  time  stood  at  740  millimetres. 
What  would  have  been  the  volume  if  observed  when  the 
barometer  stood  at  760  millimetres.     Ans.     1947  cubic  cen- 
timetres. 

(6)  A  volume  of  gas  standing  in  a  bell  glass  over  a  mer- 
cury pneumatic  trough  measured  250  cubic  centimetres.    The 
barometer  at  the  time  stood  at  754  millimetres,  and  the  level 
of  the  mercury  in  the  bell  was  found  by  measurement  to  be 
65  millimetres  above  the  surface  of  the  mercury  in  the 
trough.      Required  to  reduce  the  volume  to  the  standard 
pressure  of  760  millimetres.     Ans.     2267  millimetres. 

(6)  What  would  be  the  answer  to  the  same  problem  had 
the  trough  been  filled  with  water  ?    Ans.    246 '4  cubic  centi- 
metres. 

(7)  A  closed  vessel  which  displaces  one  litre  of  air  is 
poised  on  a  balance  with  weights  whose  volume  is  inconsid- 
erable.    The  balance  is  in  equilibrium  when  the  barometer 
stands  at  760  millimetres.    If  the  barometer  falls  to  710  milli- 
metres, how  much  weight  will  be  required  to  restore  the 
equilibrium  and  to  which  side  must  it  be  added  ?    The  tem- 
perature is  assumed  to  be  constant  at  0°.     Ans.     Weight 
required,  85  milligrammes. 

(8)  Given  the  weight  of  one  litre  of  dry  air  at  0°  C.  and 
760  millimetres,  as  above,  what  will  be  the  weight  at  0°  C. 
and  720  millimetres  ?    Ans.    It  is  obvious  that  the  weight  of 


MANOMETER.  47 

a  measured  volume  of  gas  must  be  less  in  proportion  as  the 
total  mass  of  gas  (from  which  the  measure  is  taken)  expands. 
The  formula  V  :  V  =  H' :  H  applies  to  a  limited  volume  of 
gas  under  a  variable  pressure.  Considering  now  an  invari- 
able measured  volume  of  such  a  mass  of  gas,  it  will  be  seen 
that  H  :  H'  =  W  :  W  .  ' .  760  :  720  =  1/293  :  x. 

Ex.  19.  Open  Manometer. — In  answering  the 
questions  above  the  student  will  have  learned  that 
tension  is  the  permanent  internal  elasticity  of  a 
gas,  in  virtue  of  which  it  resists  external  pressure. 
When  the  gas  is  free  to  expand,  as  when  con- 
tained by  a  bag  or  a  bell  glass  over  a  pneumatic 
trough,  that  resistance  is  the  pressure  of  the  at- 
mosphere, more  or  less  modified  by  the  medium 
through  which  it  is  transmitted ;  and,  according 
to  Mariotte's  law,  the  gas  does  expand  or  con- 
tract until  an  equilibrium  is  reached.  When  the 
gas  is  held  in  a  tight  vessel  its  volume  is  essen- 
tially invariable,  and  the  tension  is  balanced  by 
the  resistance  of  the  walls  of  the  vessel  against 
which  it  exerts,  under  ordinary  circumstances,  a 
great  pressure.  In  practical  chemistry  it  often 
becomes  a  problem  of  great  importance  to  meas- 
ure the  tension  of  a  mass  of  confined  gas,  and  the 
instrument  used  for  this  purpose  is  called  a  ma- 
nometer. One  of  the  simplest  of  these  the  stu- 
dent should  construct  and  keep  for  future  use. 

Select  a  stout  glass  tube  about  three  milli- 
metres internal  diameter  and  300  millimetres 
long ;  bend  into  the  shape  of  a  U,  and  bring  the 


48  LABORATORY  PRACTICE. 

two  arms  as  near  together  as  practicable  without 
choking  the  bend.  Leave  one  of  the  arms  straight 
open,  but  bend  the  upper  end  of  the  other  at 
right  angles  and  draw  out  to  receive  a  rubber 
connector.  Mount  on  a  wooden  stand  (easily 
made  by  tacking  together  two  pieces  of  thin 
board),  fasten  a  millimetre  scale  between  the 
tubes  and  fill  the  U  with  mercury  through  the 
open  end  until  the  level  of  the  two  columns 
stands  midway  of  the  height. 

(1)  If  on  connecting  the  manometer  with  a  glass  vessel 
the  mercury  stands  in  the  straight  arm  50  millimetres  higher 
than  in  the  other,  what  is  the  tension  of  the  gas  in  the  inte- 
rior ?  If  it  stands  50  millimetres  lower,  what  is  the  tension  ? 
Is  any  other  observation  required  in  order  to  express  the  ten- 
sion numerically  ? 

Ex.  20.  Expansion  of  Air  by  Heat.— Use  a 
plain  retort  about  fifty  cubic  centimetres  capacity. 
Dip  the  open  mouth  in  a  pan  of  water,  and  clamp 
in  position.  Heat  the  body  of  the  retort  gently 
with  a  lamp.  Allow  to  cool.  Observe  and  ex- 
plain all  the  phenomena. 

(1)  If  the  mouth  of  the  retort  were  tightly  corked  would 
the  heat  produce  any  effect  ?    How  could  this  effect  be  shown 
with  the  manometer  ?     Try  the  experiment,  first  drying  the 
retort,  tightly  corking  the  mouth,  passing  a  small  glass  tube 
through  the  cork,  and  uniting  this  tube  with  the  manometer 
by  a  stout  rubber  connector.     Try  both  heating  and  cooling 
the  retort. 

(2)  In  what  two  ways  may  the  effect  of  heat  on  a  mass  of 
gas  be  manifested  ?    Show  that  these  two  effects  must  be  pro- 
portional, and  that  one  may  be  taken  as  a  measure  of  the  other. 


LAW  OF  CHARLES  49 

Ex.  21.  Relation  of  Tension  (or  Volume)  to 
Temperature.  Law  of  Charles. — Take  a  glass 
flask  about  300  cubic  centimetres  capacity  ;  tight- 
ly cork ;  through,  the  cork  pass  two  tubes,  the 
first  a  small  tube  bent  to  connect  with  the 
manometer,  the  second  somewhat  larger  (about 
four  millimetres),  which  should  only  rise  above 
the  cork  sufficiently  to  receive  a  rubber  con- 
nector about  50  millimetres  long,  guarded  by  a 
pinch  cock.  Clamp  the  neck  to  a  firm  support, 
leaving  the  body  of  the  flask  free,  so  as  to  enable 
the  experimenter  to  bring  up  under  it  a  beaker 
sufficiently  large  to  hold  the  flask  with  space 
enough  for  the  movement  of  a  stirring  rod  all 
round  it.  Begin  by  drawing  through  the  flask 
(by  means  of  a  suction  pump)  air  dried  by  pass- 
ing through  a  chloride- of -calcium  tube.  Having 
next  connected  the  manometer  (but  leaving  the 
drying  tube  still  connected  with  the  vertical  open- 
ing), pack  the  flask  with  broken  ice.  Wait  for 
an  equilibrium  of  temperature,  and  then  close  the 
pinch  cock.  Remove  the  ice  (easily  done  by  low- 
ering the  beaker,  leaving  the  neck  clamped),  and 
replace  it  by  ice-cold  water.  Keep  stirring  the 
water  round  the  flask  until  the  temperature  has 
risen  to  5°,  a  thermometer  having  been  placed  at 
one  side,  dipping  under  the  water,  so  as  best  to 
facilitate  the  observation.  Read  now  the  differ- 
ence in  the  height  of  the  manometer  columns  and 


50  LABORATORY  PRACTICE. 

the  height  of  the  barometer.  Proceed  in  the 
same  way  until  the  temperature  has  risen  to  10°, 
and  so,  for  every  successive  five  degrees,  or  there- 
abouts, read  the  manometer,  and  also  the  barome- 
ter if  the  last  is  changing.  When  the  tempera- 
ture of  the  water  bath  nears  that  of  the  room 
apply  a  flame  to  the  beaker,  and  thus  continue 
the  observations  up  to  70°  or  80°.  Leave  the 
apparatus  to  cool  for  the  next  experiment.  Make 
now  a  table  giving  in  one  column  the  tempera- 
tures and  in  a  parallel  column  the  corresponding 
tensions,  and  seek  a  relation  between  the  values. 
They  should  correspond  to  the  proportion 

H  :  H'  =  273  +  t°  :  273  +  t'° 

Then,  since  the  volume  of  a  mass  of  gas  is,  by 
Mariotte's  law,  proportional  to  its  internal  ten- 
sion, we  have  also 

V:  V'  =  273  +  r  :273  +  f° 

(1)  Obviously  we  should  reach  the  same  result  by  meas- 
uring the  increased  volume  under  a  constant  tension  as  by 
measuring,  as  here,  the  increased  tension  under  a  constant 
volume,  but  the  last  is  the  easier  experimental  problem. 
The  proportions  which  express  the  law  of  Charles,  simply 
mean  that  the  air  increases  in  tension  or  in  volume  j|y  of 
its  tension  or  volume  at  0°.     It  is  not  necessary  however 
that  0°  should  be  taken  as  the  initial  point  of  the  actual  ex- 
periment.    It  may  be  begun  at  any  point,  and  the  tempera- 
ture raised  or  lowered  at  will,  when  the  same  relation  will 
appear.     Why  is  it  important  that  the  connecting  tube  be- 
tween flask  and  manometer  should  be  as  small  as  possible. 

(2)  At  what  temperature  would  the  tension  of  a  confined 
mass  of  at  0°  gas  be  doubled  ?    At  what  temperature  would  it 


AIR  AND  AQUEOUS  VAPOUR.  51 

theoretically  become  zero  ?  or,  assuming  the  gas  to  expand 
under  constant  pressure,  at  what  temperature  would  the  vol- 
ume be  increased  one  half  ? 

(3)  An  open  vessel  is  heated  to  819°.     What  portion  of 
the  air  which  the  vessel  contained  at  27°  remains  in  it  at 
this  temperature  ? 

(4)  Obviously  in  comparing  gas  volumes  or  gas  tensions 
we  must  have  a  standard  of  temperature,  and  0°  is  usually 
assumed  as  that  standard.     Reduce  the  following  volumes 
of  gas  measured  at  the  temperatures  and  pressures  annexed 
to  0°  and  760  millimetres. 

1.  140  c.  c.     H  =  570  mm.     t°  =  136°  '5      Ans.  70  c.  c. 

2.  320  c.  c.     H  =  950  mm.     t°  =    91°         Ans.  300  c.  c. 

3.  480  c.  c.     H  =  380  mm.     t°  =    68°  '25    Ans.  192  c.  c. 
The  following  formulae,  easily  deduced  from  the  above  pro- 
portions, will  be  useful  in  such  reductions, 


The  first  applies  when  the  volume  of  a  given  mass  of  air 
varies  under  changes  of  temperature  and  pressure,  the  sec- 
ond when  the  weight  of  air  displaced  by  a  closed  vessel,  or 
contained  in  an  open  vessel,  varies  under  the  same  circum- 
stances. The  term  to  stands  for  the  sum  (273  +  t°),  and  is 
usually  called  the  absolute  temperature.  From  the  several 
proportions  under  which  the  laws  of  Mariotte  and  of 
Charles  may  be  expressed  other  formulae  may  be  deduced 
useful  in  special  cases,  which  need  not  be  considered  here. 

Ex.  22.  Air  and  Aqueous  Vapour,  Tension  of 
Mixture.  —  Use  the  same  apparatus  as  left  from 
the  previous  experiment,  the  flask  assumed  to 
be  filled  with  dry  air  and  united  to  manometer. 
Raise  bath  to  20°,  and  maintain  steadily  at  that 
temperature.  Open  pinch  cock  until  equilibrium 
is  established,  and  then  close,  placing  the  cock  as 
low  down  on  the  rubber  tube  as  possible.  Fill 


52  LABORATORY  PRACTICE. 

the  tube  above  the  cock  with  water,  and  while 
pinching  tight  the  open  mouth  of  the  rubber  tube 
relieve  the  cock,  and,  after  squeezing  the  small  vol- 
ume of  water  into  the  flask,  again  shut.  If  all  the 
water  evaporates  add  a  further  portion  in  the 
same  way,  and  so  on.  It  will  take  some  time  be- 
fore equilibrium  is  reached.  Note  then  the  in- 
crease of  tension.  Raise  now  the  bath  to  30°,  40°, 
and  50°  successively,  and  at  each  temperature  re- 
peat the  observation,  adding  more  water  as  neces- 
sary. Remembering  now  that  the  tension  of  the 
dry  air  was  measured  by  the  height  of  the  barom- 
eter at  the  time  the  pinch  cock  was  closed,  and 
that  from  this  value  the  tension  at  any  other  tem- 
perature can  be  calculated,  assuming  also,  accord- 
ing to  a  well-known  principle,  that  when  mixed 
with  air  the  tension  of  aqueous  vapour  is  added 
to  that  of  the  air,  find  the  tension  of  aqueous 
vapour  at  each  observation,  and  compare  your  re- 
sult with  that  given  in  tables  of  the  maximum  ten- 
sion of  aqueous  vapour. 

(1)  How  can  allowance  be  made  for  a  change  in  the 
height  of  the  barometer  during  the  course  of  the  experi- 
ment ?    If  you  measure  the  tension  of  a  mass  of  air  satu- 
rated with  aqueous  vapour  at  a  given  temperature,  how  can 
you  find  what  would  have  been  its  tension  under  the  same 
condition  if  dry  ? 

(2)  A  volume  of  air  standing  in  a  graduated  tube  over  a 
water  pneumatic  trough  measures  75  cubic  centimetres.    The 
temperature  is  20°,  the  height  of  barometer  770  millimetres, 
and  the  level  of  the  water  in  the  bell  270  millimetres  above 


PREPARATION  OF  OXYGEN  GAS.  53 

that  in  the  trough.     What  would  have  been  the  volume  if 
t  =  0°,  H  =  760  millimetres,  and  the  air  dry  ? 

(3)  As  plainly  indicated  by  the  heading  of  this  division, 
we  have  in  the  last  six  experiments  studied  atmospheric 
air  as  the  type  of  aeriform  matter  and  have  tacitly  assumed 
that  all  gases  are  affected  equally  by  changes  of  pressure  or 
temperature  and  conform  sharply  to  the  two  laws  we  have 
described.  This  is  not  strictly  true,  but  the  details  of  the 
subject  would  be  out  of  place  in  an  elementary  book  and 
the  differences  are  inappreciable  in  ordinary  experiments. 
The  composition  and  chemical  relation  of  air  will  be  con- 
sidered later. 

3.  Oxygen  Gas* 

Ex.  23.  Preparation. — Take  eight  grammes  of 
pulverized  potassic  chlorate  and  two  grammes  of 
black  oxide  of  manganese.  Mix  thoroughly  in  a 
mortar  and  introduce  into  an  "  ignition  tube  "of 
Bohemian  glass.  Close  with  a  tight-fitting  cork, 
through  which  passes  a  glass  gas- delivery  tube 
leading  under  the  shelf  of  a  pneumatic  trough. 
The  pneumatic  trough  may  be  cheaply  made  of 
sheet  zinc  by  a  tin  smith,  of  such  size  as  to  hold 
one  or  two  gallons  of  water,  and  the  gas  is  best 
collected  in  ordinary  glass  fruit  jars  of  a  pint  or  a 
quart  in  capacity  which  can  be  hermetically  closed 
with  glass  covers.  The  ignition  tube  must  be 
mounted  on  some  form  of  short-tube  furnace,  and 
the  tube  heater  of  the  petroleum  stove  figured  at 
the  end  of  book  serves  an  admirable  purpose. 
Obviously,  the  cork  must  project  in  front  of  the 
furnace  sufficiently  to  prevent  scorching,  and  the 


54  LABORATORY  PRACTICE. 

powder  must  be  gathered  together  in  the  heated 
portion  of  the  tube.  The  heat  should  be  gradu- 
ally applied,  and  only  raised  to  the  highest  tem- 
perature of  the  stove  at  the  end  of  the  experi- 
ment. This  experiment  is  intended  to  give  the 
student  some  experience  with  the  methods  of  col- 
lecting and  manipulating  aeriform  matter.  He 
should  from  this  point,  if  not  before,  mount  his 
own  apparatus,  having  been  shown  how  to  fit  and 
perforate  a  cork  and  how  to  bend  a  glass  tube. 

For  the  mere  preparation  of  oxygen  gas  no 
further  details  are  required,  but  much  more  may 
be  learned  from  the  experiment. 

In  the  first  place,  the  production  of  a  gas  in- 
volves just  as  definite  a  loss  of  weight  as  would 
that  of  any  other  material.  If  several  litres  of 
oxygen  escape  from  the  ignition  tube  this  tube  will 
weigh  less  after  the  experiment  by  the  exact  weight 
of  the  aeriform  material  that  has  escaped.  To  test 
this  let  the  student  weigh  the  ignition  tube  before 
and  after  the  experiment,  and  let  him  also  measure 
the  volume  of  the  oxygen  obtained.  Under  the 
ordinary  conditions  of  a  comfortably  heated  labo- 
ratory one  litre  of  oxygen  gas  weighs  about  1*34 
gramme,  and  on  this  assumption  it  will  be  found 
that  the  total  weight  of  the  oxygen  gas  obtained 
nearly  corresponds  to  the  loss  of  weight  observed. 
There  are,  however,  obvious  causes  of  error  in 
addition  to  the  rudeness  of  the  tools  here  used 


STUDY   OF  THE   PROCESS.  55 

which  may  seriously  impair  the  sharpness  of  the 
result.  1.  Under  extreme  conditions  the  above 
assumption  may  be  materially  in  error,  although 
if  the  thermometer  and  barometer  are  observed 
correction  may  now  readily  be  made  for  this 
cause,  knowing  that  under  standard  conditions  a 
litre  of  dry  oxygen  weighs  1*4303  gramme.  2. 
The  aqueous  vapour  in  the  gas  collected  over  a 
water  pneumatic  trough  will  largely  influence  its 
volume,  and  at  a  rapidly  increasing  rate,  as  the 
temperature  rises,  and  may  produce  a  much 
greater  effect  than  has  been  allowed  in  the  esti- 
mate above  given,  although  this,  again,  may  be 
exactly  calculated.  3.  There  is  always  more  or 
less  hygrosopic  moisture  in  the  materials  used, 
which  is  expelled  with  the  gas,  and  the  error  thus 
arising  can  only  be  avoided  by  a  most  careful  pre- 
liminary drying.  Still,  as  these  errors  in  part 
compensate  each  other,  the  result  is  more  satis- 
factory than  could  be  expected. 

In  the  second  place,  the  relations  of  the  ma- 
terials in  the  combustion  tube  have  undergone  an 
essential  change.  To  test  this  shake  up  the  mate- 
rials left  in  the  tube  with  water,  filter,  and  wash 
black  residue.  Dry  and  weigh  on  filter  (the  paper 
having  been  previously  weighed).  Evaporate  the 
filtrate  to  dry  ness  in  a  porcelain  dish  (tared),  and 
weigh  the  saline  residue.  Compare  these  weights 
with  the  weights  of  oxide  of  manganese  and  of 


56  LABORATORY  PRACTICE. 

potassic  chlorate  originally  taken,  and  draw  your 
own  conclusion  as  to  the  source  of  the  oxygen  gas. 
In  the  third  place,  the  final  saline  residue  is 
not  potassic  chlorate.  Test  by  tasting,  and  also 
by  crystallizing  a  portion  of  the  residue,  and  also 
some  potassic  chlorate,  in  two  watch  glasses  side 
by  side  and  compare  form  of  crystals.  Compare 
also  black  residue  with  the  oxide  of  manganese 
used. 

(1)  What  has  happened  to  the  oxide  of  manganese  ? 

(2)  What  has  become  of  the  potassic  chlorate  ?     The 
white  residue  is  called  potassic  chloride.     What  inference 
can  you  draw  from  your  own  observation  as  to  the  differ- 
ence between  potassic  chloride  and  potassic  chlorate  ? 

(3)  Reverse  the  calculation  above.     Taking  the  loss  of 
weight  of  the  tube  as  the  weight  of  oxygen  gas  collected, 
divide  this  weight  by  the  observed  volume  and  deduce  the 
density  (the  weight  of  one  litre)  under  the  conditions  of  the 
experiment. 

Ex.  24.  (a)  Combustion. — Fill  by  displacement 
with  oxygen  gas  a  quick- sealing  (pint)  fruit  jar. 
Dry  the  jar  and  dry  the  gas  by  passing  through  a 
chloride  -  of  -  calcium  tube  the  current  from  one 
of  the  cylinders  now  familiar  in  commerce  in 
which  the  gas  is  stored  under  pressure.  Provide 
a  deflagrating  spoon  (best  made  of  sheet  iron),  so 
arranged  that  the  bowl  of  the  spoon,  hung  from  a 
cross  bar  confined  at  the  neck,  will  reach  nearly 
to  the  bottom  of  the  jar.  Line  the  bowl  with 
asbestos  paper  (previously  ignited)  and  place  on 
the  paper  less  than  half  a  gramme  of  red  phos- 


COMBUSTION.  5Y 

phorus.  Holding  the  spoon  just  over  the  mouth 
of  the  jar,  light  the  phosphorus,  and  with  a  quick 
but  deliberate  motion  plunge  the  spoon  into  the 
jar  and  seal  the  mouth.*  Notice  that  as  the  phos- 
phorus burns  away  a  white  powder  forms  in  the 
jar.  After  the  glass  is  cold,  open  the  mouth  un- 

*  It  would  be  still  better  first  to  put  the  spoon  in  place,  and, 
after  sealing  the  jar,  light  the  phosphorus  by  a  burning  glass.  This 
and  all  experiments  involving  the  deflagration  of  phosphorus  should 
be  undertaken  by  inexperienced  hands  only  under  the  most  careful 
supervision.  They  are  always  attended  with  some  risk,  and  burns 
from  phosphorus  are  painful  and  difficult  to  heal.  The  ordinary 
fruit  jar,  unless  specially  annealed,  will  not  usually  bear  the  full  in- 
tensity of  the  deflagration  without  breaking  and  sometimes  flying 
in  pieces.  The  glass  can  be  partially  annealed  by  placing  the  jar  in 
a  kettle  of  cold  water,  and,  after  bringing  the  water  to  boiling,  re- 
moving the  kettle  from  the  stove  and  leaving  it  to  cool ;  but  it  is 
also  a  safe  precaution  to  use  an  insufficient  amount  of  phosphorus 
to  exhaust  the  oxygen.  One  pint  will  consume  theoretically  0-63 
gramme  of  phosphorus,  and  if  less  than  half  a  gramme  is  used,  as 
directed,  the  deflagration  will  not  reach  a  greater  intensity  than  the 
jar  can  usually  stand.  Another  all-important  precaution  is  to  use 
red  phosphorus,  as  directed  above,  which,  although  it  does  not  ignite 
so  readily  as  ordinary  stick  phosphorus,  will  burn  very  well  if  sup- 
ported on  asbestos  paper.  In  all  cases  the  teacher  should  himself 
carefully  experiment  until  he  has  perfect  command  of  his  apparatus, 
and  then  he  can  judge  whether  this  is  a  safe  experiment  for  his 
students.  In  most  cases  he  will  undoubtedly  think  it  best  to  confine 
all  experiments  with  phosphorus  to  the  lecture  table.  Common 
phosphorus  should  never  be  used  in  the  laboratory,  and  even  with 
red  phosphorus  the  spoons  with  their  covering  when  removed  from 
the  jars  should  be  left  under  water,  and  after  any  residual  phos- 
phorus has  been  burned  off  they  should  be  scrupulously  cleaned  be- 
fore being  put  away.  Moreover,  the  room  should  be  carefully  in- 
spected after  the  class  has  left. 

N.  B. — In  all  experiments  with  fruit  jars,  in  which  it  is  impor- 
tant that  the  cover  should  hold  gas  tight,  THE  RUBBER  WASHER 

MUST  BE  WELL  GREASED. 


58  LABORATORY  PRACTICE. 

der  water,  which,  rushing  in,  will  show  that  the 
gas  has  disappeared.  This  experiment  is  the  most 
striking  illustration  we  have  that  burning  consists 
in  the  union  of  the  combustible  with  oxygen. 
The  sole  product  of  the  process  is  here  a  solid,  and 
therefore  visible. 

(5)  Combustion  in  Air. — Repeat  the  last  ex- 
periment, following  in  every  respect  the  same 
directions,  but  filling  the  jar  with  dry  air  instead 
of  oxygen.  The  action  is  less  violent,  but  the 
same  product — called  phosphoric  oxide — is  formed; 
hence  burning  in  air  is  the  same  process  as  burn- 
ing in  oxygen.  But  on  opening  the  jar  under 
water  we  find  that  only  one  fifth  of  the  volume  of 
the  air  has  been  consumed  in  the  process ;  hence 
four  fifths  of  the  volume  must  consist  of  a  differ- 
ent substance,  whose  properties  we  shall  study 
4ater. 

(c)  Combustion  of  Carbon. — Make  a  precisely 
similar  experiment  to  (a),  using  a  small  lump  of 
charcoal,  weighing  not  over  one  gramme.  In  this 
case,  although  the  coal  in  part  disappears,  there  is 
no  visible  product.  On  opening  the  mouth  of  the 
jar  it  will  be  found  to  be  still  full  of  gas.  This 
gas,  however,  is  not  oxygen.  The  properties  are 
all  changed.  Dip  a  lighted  match  into  the  jar, 
and  it  is  extinguished.  It  is  so  heavy  that  it  can 
be  poured  like  water  from  one  vessel  to  another. 
Pour  the  contents  upon  some  lime  water — or,  still 


BURNING  OF   CHARCOAL.  59 

better,  baryta  water — at  the  bottom  of  another  jar, 
and  shake  the  solution  up  with  the  gas.  Compare 
with  oxygen.  This  gas — called  carbonic  dioxide, 
often  also  carbonic-acid  gas  —  is  obviously  the 
product  of  the  burning  ;  and  this  product  was  evi- 
dently formed  by  the  union  of  carbon  and  oxygen. 
(d)  Take  a  smouldering  slow  match  and  dip 
the  lighted  end  into  a  jar  of  oxygen  ;  as  soon  as 
the  match  inflames  remove  and  repeat  the  trial, 
thus  following  the  level  of  the  gas  as  it  is  con- 
sumed. 

(1)  Do  the  facts  which  you  have  observed  in  (a)  justify 
the  conclusion  that  the  burning  consisted  in  a  union  of 
phosphorus  and  oxygen  ?  and  that  the  white  residue  was 
the  product  of  that  union  ? 

(2)  Do  the  further  facts  observed  in  (6)  justify  the  con- 
clusion that  one  fifth  of  the  volume  of  atmospheric  air  con- 
sists of  oxygen  ? 

(3)  What  becomes  of  the  charcoal  in  (c)  ?  and  what  must 
be  the  composition  of  the  aeriform  product  left  in  the  jar  ? 
Show  that  the  facts  observed  support  your  inference. 

(4)  So  far  as  your  own  observation  has  extended,  what 
are  the  properties  of  oxygen  gas  ?    Is  it  the  same  material 
which  was  obtained  as  one  of  the  products  in  the  decomposi- 
tion of  water  ?    How  would  you  recognize  the  gas  if  you 
met  with  it  as  the  product  of  a  chemical  process  ? 

(5)  What  is  the  nature  of  combustion  so  far  as  illus- 
trated by  the  above  experiments  ? 

4.  Hydrogen  Gas. 

Ex.  25.  Preparation. — Prepare  hydrogen  gas 
by  pouring  a  mixture  of  one  part  of  sulphuric  acid 
and  five  parts  of  water  over  clippings  of  sheet 


60  LABORATORY  PRACTICE. 

zinc  in  a  glass  flask.  Connect  the  glass  flask  by 
means  of  a  cork  and  glass  tube  with  a  pneumatic 
trough,  and  collect  the  gas  in  glass  fruit  jars  after 
the  air  has  been  driven  from  the  flask.  As  each 
jar  fills  it  may  be  lifted  from  the  water  and  the 
gas  burned  at  the  mouth.  For  other  experiments 
a  self-regulating  generator  should  be  provided, 
from  which  the  gas  can  be  drawn  as  required.* 

Ex.  26.  Density  of  Hydrogen.  —  The  great 
lightness  of  hydrogen  gas  can  be  illustrated  by 
the  teacher  in  various  ways — as  by  filling  a  small 
balloon,  or  blowing  soap-bubbles  with  it ;  also 
by  decanting  it  from  a  large  jar  to  a  smaller,  hold- 
ing both  jars  with  the  mouths  down.  The  student 
may  test  this  very  striking  quality  for  himself  by 
filling  a  jar  with  the  gas  by  displacement,  only 

*  Such  a  generator  can  easily  be  made  from  old  glass  bottles, 
and  will  be  useful  in  several  experiments.  For  the  inside  bell  use  a 
half-pint  tincture,  of  thin  glass.  Through  the  bottom  of  such  a 
bottle  bore  several  holes  three  to  four  millimetres  in  diameter,  using 
as  a  drill  the  sharpened  point  of  a  three-cornered  file  dipped  in 
turpentine  or  kerosene.  Fill  the  bottle  with  zinc  clippings,  tightly 
cork,  and  pass  through  the  cork  a  gas-evolution  tube  guarded  by  a 
pinch  cock.  Clamp  this  bell  firmly  by  means  of  wooden  stays  fast- 
ened by  twine,  in  a  considerably  larger,  open,  tumbler-shaped  vessel, 
which  can  be  made  by  cutting  off  (with  a  hot  coal,  "  sprengkohlen  ") 
the  neck  of  a  common  quart  wine  bottle,  or  a  tall  glass  beaker  may 
be  used  for  the  purpose.  Lastly,  fill  the  open  vessel  with  dilute  sul- 
phuric acid.  On  opening  the  cock  the  acid  will  flow  into  the  bell, 
and  the  gas  generated  by  its  action  on  the  zinc  will  soon  drive  out 
the  air.  On  then  closing  the  cock  the  bell  will  fill  with  gas  and 
drive  back  the  acid  water,  when  the  chemical  action  will  cease  and 
the  apparatus  be  left  in  a  condition  to  yield  a  constant  supply  of 
hydrogen. 


BURNING  OF  HYDROGEN  GAS.  61 

displacing  the  air  from  above  downwards  instead 
of  the  reverse,  as  in  the  case  of  oxygen.  This  is 
most  readily  done  by  resting  the  open  mouth  of 
the  jar  on  a  square  of  cardboard,  supported  by  the 
ring  of  a  retort  stand,  and  passing  the  glass  deliv- 
ery tube  from  the  generator  through  a  hole  in  the 
cardboard  (which  it  should  tightly  fit)  to  the  very 
top  of  the  jar.  When  the  jar  is  full,  and  while 
the  gas  current  is  still  flowing,  the  tube  should  be 
slowly  withdrawn  ;  and  a  little  address  is  required 
to  slip  under  the  cap  at  the  right  moment,  so  as  to 
prevent  air  from  mixing  with  the  hydrogen  at  the 
open  mouth. 

Hydrogen  gas  is  about  fourteen  and  a  half 
times  lighter  than  air,  about  sixteen  times  lighter 
than  oxygen,  and,  under  standard  conditions,  one 
litre  of  hydrogen  weighs  very  closely  0*09  gramme. 

Ex.  27.  Combustion  of  Hydrogen. — The  pro- 
duction of  water  by  the  burning  of  hydrogen 
should  be  shown  on  as  large  a  scale  as  possible  by 
the  teacher,  and  the  apparatus  described  in  the 
author's  New  Chemistry  is  admirably  adapted  to 
this  purpose.  To  observe  the  same  effect,  let  the 
student  replace  the  delivery  tube  of  the  flask  used 
in  Ex.  25  with  a  short  piece  of  tube  drawn  out  to 
a  jet ;  and  after  making  assurance  doubly  sure 
that  all  the  air  has  been  driven  from  the  flask,* 

*  The  explosion  of  "  hydrogen  flasks  "  is  a  very  frequent  acci- 
dent in  chemical  laboratories.     In  such  cases  the  injury  is  usually 
5 


62  LABORATORY  PRACTICE. 

light  and  burn  the  hydrogen  at  the  jet.  Hold 
now  a  cold  and  dry  jar  over  the  jet  until  the  moist- 
ure which  condenses  collects  in  drops  to  a  suffi- 
cient extent  to  render  the  nature  of  the  product 
evident.  Test  in  a  similar  way  the  products  of  a 
flame  of  common  illuminating  gas ;  and,  after  a 
perceptible  amount  of  water  has  been  formed, 
shake  up  in  the  jar  some  lime  water,  and  prove  by 
the  resulting  fcurbidity  that  carbonic  dioxide  has 
also  been  formed.  Test  likewise  the  products  of 
the  flame  of  a  candle.  In  this  connection  the 
teacher  should  discuss  at  length  the  general  feat- 
ures of  the  burning  of  hydrocarbon  fuels,  includ- 
ing both  the  nature  of  the  products  and  the  stages 
of  the  process. 

Ex.  28.  Nature  of  Flame.— Make  the  follow- 
ing experiments  with  the  flame  of  a  Bunsen  lamp 
with  the  air  valve  shut,  and  also  as  far  as  prac- 
ticable with  the  flame  of  a  candle.  1.  Press  the 
flame  down  with  a  broken  bit  of  glass,  and  notice 
that  the  shell,  or  mantle,  only  is  luminous.  2. 
Adjust  a  piece  of  glass  tube  so  as  to  draw  off  the 
combustible  gas  which  forms  the  cone  of  the 
flame,  and  notice  that  it  may  be  lighted  at  the 
upper  end  of  the  tube.  3.  Press  down  on  the 

caused  by  the  scattering  of  the  glass ;  and  the  danger  can  be  in 
great  measure  prevented  by  wrapping  the  flask  in  a  towel  before 
lighting  the  jet.  For  this  experiment  it  is  safer  to  us&  a  hydrogen 
generator  that  has  been  tested. 


BURNING.  63 

flame  a  screen  of  fine- wire  ganze,  and  notice  that, 
although  the  combustible  gas  passes  through,  the 
burning  mantle  is  cut  off.  4.  Press  down  on  the 
flame  the  back  of  a  cold  iron  spoon,  and  notice  the 
soot  which  collects  upon  it.  5.  Open  the  air  valve 
of  the  Bunsen  burner,  and  notice  that  when  the 
flame  loses  its  luminous  power  no  soot  is  deposited. 
By  means  of  these  and  similar  experiments  the 
teacher  should  enforce  the  following  conclusions  : 
1.  That  flame  is  always  burning  gas.  2.  That  the 
action  takes  place  in  the  outer  mantle  of  the 
flame.  3.  That  a  combustible  will  not  take  fire 
and  continue  burning  unless  its  temperature  is 
raised  to  the  point  of  ignition  and  maintained  at 
that  temperature.  4.  That  in  ordinary  burning 
the  required  temperature  is  maintained  by  the 
heat  developed  from  the  union  of  the  combustible 
with  the  oxygen  of  the  air.  5.  That  when  the 
burning  gas  of  a  flame  is  cooled  below  its  point  of 
ignition  the  flame  goes  out,  and  hence  the  efficacy 
of  the  wire  gauze  in  the  safety  lamp.  6.  That 
charcoal,  not  being  volatile,  burns  without  flame, 
but  that  here  also,  if  the  glowing  coals  are  cooled 
below  a  red  heat,  the  union  with  oxygen  stops, 
and  the  fire  is  extinguished.  In  this  connection 
the  student  should  be  shown  the  use  of  the  mouth 
blowpipe,  and  the  effects  of  the  reducing  and  oxi- 
dizing flames  should  be  explained. 


64  LABORATORY  PRACTICE. 

(1)  It  is  not  expected  that  the  student  will  understand 
the  nature  of  the  chemical  process  in  Ex.   25.     This  will 
hereafter  appear  ;  but  he  should  clearly  recognize  at  this 
stage  that  the  product  is  the  same  combustible  gas  obtained 
in  the  decomposition  of  water. 

(2)  Show  that  Ex.   27  confirms  the  analysis  of  water 
in  Ex.  16. 

(3)  From  what  does  the  carbonic  dioxide  formed  by  a 
burning  candle  come  ?    In  what  respects  does  the  flame  of 
a  candle  differ  from  the  flame  of  hydrogen  ?    What  is  the 
cause  of  the  differences  ? 


5.  Sulphur. 

Ex.  29.  Specific  Characters.  —  The  student 
should  be  given  a  roll  of  brimstone  and  asked 
to  study  and  describe  its  distinguishing  proper- 
ties, including  colour,  hardness,  tenacity,  specific 
gravity,  fusibility,  volatility,  colour  of  vapour, 
and  solubility  in  ordinary  solvents. 

Ex.  30.  Melting  and  Boiling  Points.—  The 
melting  point  may  be  determined  by  heating  a 
small  bit  of  sulphur  in  a  glass  quill  tube  closed 
at  the  lower  end  by  means  of  a  bath  of  some 
liquid  in  which  the  tube  is  dipped.  The  bath 
generally  used  for  determining  melting  points  is 
a  small  beaker  glass  containing  sulphuric  acid 
heated  over  a  lamp  ;  but,  as  the  use  of  hot  sul- 
phuric acid  is  not  unattended  with  danger  in  the 
hands  of  the  inexperienced,  melted  paraffine  or 
castor  oil  had  better  be  substituted  in  this  experi- 
ment. The  temperature  of  the  bath  at  which  the 


SULPHUR.  65 

sulphur  begins  to  melt  is  observed  by  means  of 
a  thermometer  hanging  so  that  its  bulb  dips  in 
the  same  bath  at  the  side  of  the  tube ;  but  the 
student  will  require  some  instruction  and  practice 
before  he  can  obtain  accurate  results  with  this  ap- 
paratus. To  measure  directly  the  boiling  point 
of  sulphur  we  require  a  peculiar  thermometer 
made  by  Geissler,  which  is  filled  under  pressure, 
and  indicates  temperature  up  to  450°.  The  sul- 
phur is  boiled  in  a  small  flask  with  a  long  neck, 
and  the  thermometer  should  be  suspended  so  that 
its  bulb  hangs  in  the  midst  of  the  deep-red  va- 
pour. Such  thermometers  cost  in  Europe  three 
dollars  each,  and  unless  provided  this  observation 
must  be  omitted. 

Ex.  31.  Modifications  of  Sulphur. — First,  dis- 
solve a  gramme  of  sulphur  in  sulphide  of  carbon, 
and  allow  the  solution  to  evaporate  spontane- 
ously, when  the  sulphur  will  crystallize.  (Sul- 
phide of  carbon  is  very  volatile  and  inflamma- 
ble. HAVE  NO  LIGHTS  NEAR  BY.)  Second,  melt 
enough  sulphur  to  nearly  fill  a  small  beaker,  tak- 
ing care  that  the  temperature  does  not  rise  much 
above  the  melting  point.  Let  the  vessel  cool  un- 
til a  crust  begins  to  form,  and  then  promptly 
pour  out  what  remains  liquid,  which  will  leave 
the  beaker  lined  with  crystals  of  sulphur  having 
a  very  different  form  and  color  from  those  first 
obtained.  Third,  melt  some  sulphur  in  a  test 


66  LABORATORY  PRACTICE. 

tube,  and  raise  the  temperature  until  the  liquid, 
at  first  very  limpid,  becomes  thick  and  pasty,  and 
then  pour  the  material  out  in  a  fine  stream  into 
a  basin  of  cold  water.  Let  now  the  student  study 
and  describe,  as  well  as  he  can,  the  differences  be- 
tween these  three  conditions  of  sulphur. 

(1)  Are  the  several  modifications  of  sulphur  the  same 
substance  or  different  substances  ? 

Ex.  32.  Combustion  of  Sulphur.  —  Burn  a 
small  amount  of  sulphur  in  a  jar  of  oxygen  gas, 
making  the  experiment  as  described  in  Ex.  24. 
The  action  is  far  less  violent  than  in  that  experi- 
ment ;  and,  in  order  to  facilitate  the  burning,  it  is 
better  to  use  shreds  of  asbestos  soaked  with  sul- 
phur (when  melted)  instead  of  a  lump  of  brim- 
stone. The  asbestos  is  not  acted  on,  and  prevents 
the  sulphur,  when  melted  by  the  heat,  from  run- 
ning together.  When  the  jar  is  opened  it  will  be 
found  that  it  is  still  full  of  an  aeriform  material, 
but  that  the  new  product,  although  having  the 
same  volume,  is  wholly  different  from  the  oxygen 
gas  with  which  the  experiment  began.  The  prod- 
uct is  obviously  a  compound  of  sulphur  and  oxy- 
gen, and  must  weigh  more  than  the  initial  oxygen 
by  just  the  weight  of  sulphur  burned.  Indeed, 
as  will  hereafter  appear,  it  weighs  just  twice  as 
much  as  the  same  volume  of  oxygen.  Again,  un- 
like oxygen,  it  has  a  very  suffocating  odor,  and 


SULPHURIC  ACID.  67 

by  this  will  be  recognized  as  the  same  product 
which  is  so  offensive  from  a  burning  match.  Let 
the  student  immerse  in  the  gas  a  red  rose,  or  some 
other  highly  coloured  flower,  and  witness  the  re- 
markable bleaching  power.  Let  him  then  dis- 
solve the  rest  of  the  gas  by  shaking  up  50  cubic 
centimetres  of  water  in  the  jar.  The  gas  is  called 
sulphurous  oxide,  and  the  solution  in  water, 
which  has  acid  qualities,  is  well  known  as  sul- 
phurous acid.  Taste  the  solution.  Dip  into  it, 
momentarily,  a  strip  of  paper  coloured  blue  with 
litmus. 

(1)  Justify  the  inference  that  sulphurous  oxide  is  com- 
posed of  sulphur  and  oxygen. 

Ex.  33.  Production  of  Sulphuric  Oxide.  — 
Take  a  short  length  of  small  combustion  tubing 
(100  millimetres  long  and  from  6  to  8  millimetres 
bore),  fill,  but  not  too  tightly,  the  middle  por- 
tion of  the  tube  with  platinized  asbestos,*  tightly 
cork  one  end  and  draw  out  the  other  to  a  small 
tubulature ;  through  a  perforation  in  the  cork 
pass  a  small  glass  tube.  Wire  the  combustion 
tube  in  a  horizontal  position  to  the  ring  of  a  re- 
tort stand.  Take  also  a  test  tube  corked  tight- 

*  Platinized  asbestos  can  be  purchased  of  dealers  in  chemicals 
but  is  easily  made  by  drenching  asbestos  wool  with  solution  of  plati- 
num chloride  and  then  igniting.  Asbestos  paper  can  be  treated  in 
this  way,  and  affords  a  convenient  preparation,  since  it  can  be  rolled 
into  a  shape  that  will  just  fit  and  fill  the  tube. 


68  LABORATORY  PRACTICE. 

ly ;  make  two  perforations  in  the  cork.  Through 
one  of  these  pass  an  inlet  tube  extending  to  the 
bottom  of  the  test  tube,  and  through  the  other  an 
outlet  tube  only  extending  through  the  cork. 
Stand  the  test  tube  in  a  beaker  and  pack  round 
it  broken  ice  and  connect  the  inlet  tube  by  a  rub- 
ber connector  with  the  tubulature  of  the  combus- 
tion tube.  Connect  the  other  end  of  the  combus- 
tion tube  with  a  bent  glass  tube  that  will  reach  to 
the  bottom  of  a  fruit  jar.  Use  a  quart  fruit  jar, 
and,  having  filled  it  by  displacement  with  dry 
oxygen — Ex.  24  (a) — repeat  Ex.  32,  using  only 
half  a  gramme  of  sulphur.  This  amount  is  not 
sufficient  to  exhaust  the  oxygen,  so  that  at  the 
end  of  the  combustion  there  must  remain  in  the  jar 
a  mixture  of  sulphurous  oxide  and  oxygen  gases. 
After  the  jar  has  cooled  uncover,  remove  the  defla- 
grating spoon,  and  adjust  the  bent  tube  so  that 
it  leads  to  the  bottom  of  the  jar,  as  above  de- 
scribed. Place  a  Bunsen  lamp  under  the  combus- 
tion tube  and  heat  the  platinized  asbestos  to  dull- 
red  heat.  Connect,  lastly,  the  other  end  of  the 
apparatus  with  an  aspirator,  and  slowly  draw  the 
contents  of  the  jar  over  the  asbestos  and  through 
the  test  tube.  Notice  that  during  the  process  the 
platinized  asbestos  undergoes  no  change  what- 
ever, but  there  will  collect  in  the  test  tube  a  con- 
siderable amount  of  a  white  crystalline  solid. 
After  dismounting  the  apparatus  dissolve  the 


SULPHURIC  ACID.  69 

substance  collected  in  the  cold  test  tube  in  a  few 
cubic  centimetres  of  water,  and  compare  the  prop- 
erties of  this  solution  with  those  of  the  dilute  sul- 
phuric acid  used  in  Ex.  25.  For  this  purpose 
prepare  a  very  weak  solution  of  baric  chloride 
and  half  fill  with  it  two  test  tubes  standing  side 
by  side.  Add  to  each  a  few  drops  of  hydro- 
chloric acid.  Add  to  the  first  a  few  drops  of 
what  is  known  to  be  diluted  sulphuric  acid,  such 
as  used  in  Ex.  25  ;  add  to  the  second  the  product 
of  the  last  experiment,  and  compare  results.  Re- 
peat now  Ex.  24  (a),  and  dissolve  the  white  prod- 
uct of  that  combustion  in  a  few  cubic  centimetres 
of  water ;  notice  the  time  taken  for  complete  so- 
lution. Taste,  test  with  litmus  paper,  and  also 
with  baric  chloride  in  the  same  way  as  above. 
Obviously,  then,  both  phosphorus  and  sulphur, 
by  uniting  with  oxygen  in  the  process  of  combus- 
tion, yield  compounds  which,  when  dissolved  in 
water,  give  products  that  have  a  strong  acid  taste. 

Phosphorus,  oxygen,  and  water      yield  phosphoric  acid. 
Sulphur,  oxygen,  and  water  "     sulphurous     " 

"        more  oxygen,  and  water      "     sulphuric       " 

How  far  these  cases  are  illustrations  of  a  general 
principle  will  appear  hereafter. 

(1)  Are  we  justified  in  drawing  the  conclusion  from  this 
experiment  that  sulphuric  oxide  differs  from  sulphurous 
oxide  and  sulphuric  acid  from  sulphurous  acid  only  in  hold- 
ing more  oxygen  in  combination  ? 


70  LABORATORY  PRACTICE. 

(2)  On  the  basis  of  the  facts  hitherto  observed,  is  the  evi- 
dence above  given  (that  the  final  product  of  this  experiment 
is  identical  with  common  sulphuric  acid)  satisfactory  ? 


6.  Chlorine. 

Ex.  34.  Preparation  of  Hydrochloric- Acid 
Gas. — Mix  20  grammes  of  sulphuric  acid  with  4 
grammes  of  water,  pouring  the  acid  slowly  into 
the  water,  and  when  cold  add  this  mixture  to  10 
grammes  of  powdered  common  salt  in  a  glass 
flask.  Cork  the  flask  and  provide  a  glass  tube 
passing  through  the  cork  and  leading  to  the  bot- 
tom of  a  fruit  jar.  Cover  the  mouth  of  the  jar 
with  a  card  (Ex.  26),  and  collect  the  heavy,  suffo- 
cating gas  by  displacement.  Three  full  jars  will 
be  required  for  further  experiments.  These,  be- 
fore filling,  must  be  thoroughly  dried,  and  the 
dense  fumes  which  form  when  hydrochloric-acid 
gas  mixes  with  the  air  will  always  show  when  a 
jar  is  filled  to  overflowing.  To  one  of  the  jars 
prepared  as  above  add  10  grammes  of  water, 
which  will  instantly  absorb  the  contents  ;  and  the 
solution  thus  obtained,  although  weaker,  is  the 
same  preparation  as  the  liquid  hydrochloric  acid 
so  much  used  in  chemical  laboratories.  The  com- 
mercial acid  often  contains  over  four  hundred 
times  its  volume  of  dissolved  gas. 

Ex.  35.  Composition  of  Hydrochloric- Acid 
Gas. — To  another  jar  of  the  gas  obtained  in  the 


CHLORINE  GAS.  71 

last  experiment  add  20  or  30  grammes  of  sodium 
amalgam,  and,  instantly  sealing  the  jar,  shake  up 
the  amalgam  with  the  gas  as  long  as  absorption 
continues.  Then  open  the  jar  under  water,*  which 
will  rush  in  and  show  that  one  half  of  the  volume 
has  been  absorbed.  Apply  a  lighted  match  to  the 
residual  gas,  and  it  will  be  found  to  be  hydrogen. 
The  only  other  ingredient  of  hydrochloric-acid  gas 
will  appear  in  the  next  experiment. 

Ex.  36.  Preparation  of  Chlorine  Gas. — Fill  a 
flask  (50  cubic  centimetres  capacity),  fitted  with  a 
perforated  rubber  stopper  and  outlet  tube,  to 
three  fourths  of  its  capacity  with  lumps  of  black 
oxide  of  manganese.  Pour  upon  these  lumps 
strong  liquid  hydrochloric  acid  so  as  to  fill  the 
interstices  only.  Allow  the  flask  to  stand  for  a 
short  time,  and  then  apply  a  gentle  heat.  A  yel- 
lowish gas  comes  off  in  abundance,  which  is  much 
heavier  than  the  air,  and  can  be  collected  by  dis- 
placement if  the  outlet  tube  reaches  quite  to  the 
bottom  of  the  jars  used  for  the  purpose,  and  the 
mouth  is  covered  with  a  disk  of  cardboard,  as 
already  described.  Chlorine  gas  is  very  suffocat- 
ing, and  the  smallest  puff,  if  inhaled,  may  pro- 
duce serious  results.  This  experiment  should 
therefore  be  performed  with  extreme  caution, 

*  Use  a  glass  or  porcelain  dish  to  hold  the  water.  Not  the 
pneumatic  trough,  which  would  be  corroded  by  the  mercury  falling 
into  it. 


72  LABORATORY  PRACTICE. 

either  in  the  open  air  or  under  a  hood  with  a 
strong  draught.  The  colour  of  the  gas  shows  when 
a  jar  is  full,  and  three  jars  thus  filled  are  required 
for  the  following  experiments :  1.  In  the  first  of 
these  jars  plunge  some  tinsel  or  other  metal  leaf 
hanging  from  the  end  of  a  long  stick,  and  almost 
every  metal  will  at  once  enter  into  direct  union 
with  the  chlorine,  often  with  ignition.  2.  Place 
mouth  to  mouth  a  jar  of  hydrogen  over  a  jar  of 
chlorine,  and,  holding  the  open  mouths  together 
confined  by  a  rubber  band,  invert  the  two,  and  in 
a  few  moments,  when  the  gases  have  mixed, 
loosely  cover  (but  on  no  account  seal)  both  of  the 
jars.  If  now  one  of  the  jars  is  exposed  to  bright 
daylight  (not  direct  sunlight),  a  gradual  union  be- 
tween the  chlorine  and  the  hydrogen  gases  will 
take  place,  and  after  the  yellow  colour  has  disap- 
peared the  product  can  easily  be  recognized  as  hy- 
drochloric acid  gas.  The  second  jar,  kept  in  the 
dark,  will  undergo  no  change,  and  can  be  used  for 
comparison.  If  now  this  jar  is  exposed  to  direct 
sunlight,  the  same  combination  will  suddenly  take 
place  with  explosive  violence.*  This  experiment 
is  a  dangerous  one,  and  should  only  be  made  with 
the  greatest  caution,  the  mouth  of  the  jar  being 

*  There  is  more  or  less  danger  in  all  stages  of  this  experiment, 
not  only  from  the  violence  of  the  explosion,  but  also  from  the  risk  of 
breathing  a  puff  of  chlorine.  It  should  never  be  intrusted  to  care- 
less hands,  nor  indeed  to  the  hands  of  any  student  before  all  the 
necessary  precautions  have  been  pointed  out  and  enforced. 


CARBON.  Y3 

loosely  covered  with  a  pasteboard  disk.  The  com- 
position of  hydrochloric  acid  is  thus  fully  estab- 
lished. 3.  Into  the  third  jar  of  chlorine,  prepared 
as  above,  pour  200  cubic  centimetres  of  water, 
and,  after  closing  the  jar,  shake  the  water  up  with 
the  gas,  which  will  be  almost  completely  absorbed 
and  impart  its  colour  to  the  solution.  Soak  then 
in  the  water  a  strip  of  calico  printed  with  madder, 
and  notice  how  rapidly  the  colour  is  discharged. 
In  this  connection  the  use  of  chlorine  as  a  bleach- 
ing agent  should  be  explained. 

(1)  What  proof  has  been  given  that  hydrochloric  acid 
gas  consists  solely  of  hydrogen  and  chlorine  ?    Do  our  ex- 
periments show  in  what  proportions  hydrogen  gas  and  chlo- 
rine gas  combine  by  volume  ? 

(2)  What  is  the  composition  of  liquid  hydrochloric  acid? 
Have  you  noticed  any   difference   between   the  union  of 
hydrochloric-acid  gas  with   water  and  ordinary  solution  ? 
Does  the  composition  of  liquid  hydrochloric  acid  conform  to 
the  general  scheme  exhibited  under  Ex.   33  ?    Point  out 
agreements  and  differences. 

(3)  Compare   sulphuric    acid  and    liquid  hydrochloric 
acid,  using  the  ordinary  laboratory  acids  diluted  with  four 
or  five  times  their  volume  of  water.     Try  taste,  litmus  paper, 
and  action  on  zinc  clippings.     Also  evaporate  a  few  drops 
of  each  on  watch  glasses  and  observe  effects. 

7.  Carbon. 

Ex.  37.  (a)  Preparation  of  Charcoal. — Take 
small  billets  of  three  or  four  different  kinds  of 
wood,  including  the  densest  and  lightest  that  can 
be  procured.  Cover  with  sand  in  an  iron  crucible. 


74  LABORATORY  PRACTICE. 

and  heat  to  redness  until  the  smoking  stops.  It 
is  best  to  light  the  gas  thus  given  off  above  the 
sand  to  prevent  it  from  escaping  into  the  room. 

(b)  Heat  to  redness  over  a  lamp  in  a  porcelain 
crucible  some  lumps  of  sugar  as  long  as  any  vapor 
is  evolved.  In  this  connection  the  general  facts  in 
regard  to  the  composition  of  organic  substances 
should  be  briefly  stated,  and  the  relations  of 
carbon  as  the  non-volatile  skeleton  of  organized 
matter  should  be  explained.  Also,  the  relations 
of  diamond  and  graphite  to  charcoal  should  be 
stated,  and  compared  with  those  between  the  dif- 
erent  states  of  sulphur. 

Ex.  38.  Specific  Characters  of  Charcoal.— The 
student  should  be  asked  to  study  the  distinguish- 
ing characters  of  the  charcoal  prepared  in  the  last 
experiment,  and  to  compare  these  with  those  of 
sulphur  already  studied.  If  the  work  is  judicious- 
ly directed  and  criticised  this  exercise  will  be  very 
instructive  ;  and,  for  the  very  reason  that  the  two 
substances  are  so  different,  it  will  be  a  good  prepa- 
ration for  comparing  hereafter  two  substances 
which  are  closely  alike.  As  charcoal  is  a  porous 
body  whose  external  volume  depends  on  that  of 
the  organized  material  from  which  it  was  made, 
the  density  must  obviously  be  left  out  of  consid- 
eration in  this  comparison. 

Ex.  39.  (a)  Preparation  of  Carbonic  Dioxide 
(Carbonic- Acid  Gas). — This  gas,  as  we  have  seen, 


CARBONIC  DIOXIDE.  75 

is  formed  by  the  burning  of  charcoal  (Ex.  24,  c) ;  it 
is  also  easily  made  by  the  action  of  liquid  hydro- 
chloric acid  on  marble  (calcic  carbonate).  Half 
fill  a  glass  flask  (250  cubic  centimetres  capacity) 
with  small  lumps  of  marble.  Pour  upon  the  mar- 
ble common  muriatic  acid  (the  commercial  name 
of  crude  liquid  hydrochloric  acid)  mixed  with 
three  times  its  volume  of  water.  Connect  with  a 
pneumatic  trough  and  collect  in  the  usual  way. 
Half  fill  one  of  the  jars,  and,  after  closing,  shake 
the  gas  up  with  the  remaining  water.  Open  from 
time  to  time  to  admit  air  until  the  absorption 
ceases.  At  the  ordinary  pressure  of  the  air  water 
will  absorb  its  own  volume  of  carbonic-dioxide 
gas,  and  soda  water  is  the  same  solution  under 
pressure.  This  aqueous  solution  may  be  called 
carbonic  acid.  Dip  into  it  a  strip  of  blue  litmus 
paper  and  notice  the  difference  between  the  effect 
of  this  weak  volatile  acid  and  that  of  a  strong 
fixed  acid  like  sulphuric  acid  when  both  are 
equally  dilute. 

(b)  What  was  the  Source  of  Carbonic  Dioxide 
in  Last  Experiment  f — To  answer  this  question, 
prepare  some  lime  water  from  quicklime  slaked  as 
in  Ex.  15.  Fill  a  quart  fruit  jar  somewhat  over 
one  half  (four  sevenths)  with  lime  water  and  the 
rest  with  carbonic  dioxide  (pouring  in  the  heavy 
gas  just  as  you  would  a  liquid).  Close,  and  shake 
the  gas  and  water  well  together,  admitting  air  from 


76  LABORATORY  PRACTICE. 

time  to  time  as  the  absorption  goes  on.  Allow 
the  precipitate  to  settle  ;  pour  off  the  clear  water, 
collect  the  precipitate  on  a  filter,  wash,  and  dry. 
Transfer  next  the  powder  to  a  test  tube  and  pour 
on  a  few  cubic  centimetres  of  dilute  hydrochloric 
acid.  Notice  the  effervescence  and  recognize  as 
carbonic  dioxide  the  gas  evolved.  To  make  the 
demonstration  complete  the  student  must  be  told 
that  marble  is  one  of  the  many  mineral  forms  of 
carbonate  of  lime  and  is  chemically  the  same 
substance  as  the  white  powder  thus  prepared. 
The  proof  of  this  will  appear  later.  In  this  con- 
nection the  important  relations  of  carbonic  di- 
oxide in  nature  should  be  discussed,  its  pres- 
ence in  the  breath  should  be  shown,  and  its 
association  with  alcohol  as  a  product  of  fermen- 
tation and  its  presence  in  beer,  sparkling  wine, 
and  other  effervescing  drinks,  should  be  ex- 
plained. 

(c)  Repeat  Ex.  24  (c),  and  after  the  jar  has 
cooled  open  the  mouth  under  water.  There  will 
be  no  expansion  of  the  aeriform  product,  and  no 
contraction  except  that  due  to  the  slow  solution 
of  the  gas  in  water.  Obviously,  then,  a  given 
volume  of  oxygen  gas  yields  the  same  volume  of 
carbonic-dioxide  gas,  and  the  last  must  weigh 
more  than  the  first  by  the  weight  of  the  charcoal 
which  the  oxygen  gas  absorbs  in  the  process  of 
burning, 


CARBONIC   OXIDE.  77 

(1)  Is  the  volume  of  the  atmosphere  altered  by  the 
smoke  which  our  fires  pour  into  it  ? 

Ex.  40.  (a)  Production  of  Carbonic  Oxide. — 
Provide  two  rubber  gas  bags,  holding  about  one 
litre  each.  Connect  these  with  the  ends  of  a 
length  of  hard  Bohemian  glass  tubing,  which 
should  be  filled,  but  not  tightly  packed,  with  fine- 
ly pulverized  charcoal  that  has  been  thoroughly 
burned.  Having  filled  one  of  the  bags  to  about 
one  half  of  its  capacity  with  carbonic  dioxide, 
heat  the  tube  over  two  or  more  Bunsen  lamps 
(best  a  gas  tube  furnace)  to  a  full  red  heat,*  and 
pass  the  gas  slowly  backwards  and  forwards  so 
long  as  any  increase  of  volume  is  perceptible,  and 
at  last  it  will  be  found  that  the  volume  has 
doubled.  Remove  now  the  full  bag  and  transfer 
a  portion  of  the  product  to  a  small  glass  jar  over 
the  pneumatic  trough.  Lift  the  jar  from  the 
water  with  the  mouth  down  and  the  gas  will  not 
at  once  escape,  because  the  new  product  is  even 
lighter  than  air.  It  may  now  be  lighted  at  the 
mouth  of  the  jar  and  the  peculiar  color  of  the 
flame  noticed  and  the  product  of  the  combustion 
shown  to  be  carbonic  dioxide.  In  this  experi- 
ment it  is  evident  that  the  carbonic  dioxide  must 


*  The  kerosene  stove  does  not  give  a  sufficiently  high  tempera- 
ture. By  burning  alcohol  in  the  stove,  however,  the  requisite  heat 
can  be  obtained ;  but  care  should  be  taken  to  avoid  explosions,  and 
it  will  be  safer  to  reserve  this  experiment  for  the  lecture  table. 

6 


78  LABORATORY  PRACTICE. 

have  united  with  the  material  of  the  charcoal,  and 
the  new  product,  called  carbonic  oxide,  must  dif- 
fer from  carbonic  dioxide  only  in  containing  more 
carbon  or,  what  amounts  to  the  same  thing,  pro- 
portionally less  oxygen.  The  same  inference  may 
be  drawn  from  the  fact  that  in  burning  carbonic 
oxide  changes  back  to  carbonic  dioxide. 

(b)  Add  ten  grammes  of  oxalic  acid  to  a  small 
flask,  corked  and  fitted  with  an  evolution  tube 
leading  to  a  pneumatic  trough.  Pour  over  this, 
crystalline  solid  five  or  six  times  its  weight  of 
strong  sulphuric  acid  (oil  of  vitriol).  Support  (on 
a  retort  stand)  the  flask,  protected  by  a  square  of 
asbestos  paper,  and  apply  gentle  heat.  The  gas 
which  comes  over  copiously  is  a  mixture  of  car- 
bonic oxide  and  carbonic  dioxide.  The  last  will 
be  slowly  absorbed  by  the  water,  and  very  rapidly 
if  some  caustic  soda  is  added  to  the  pneumatic 
trough.  Collect  in  fruit  jar,  and,  after  allowing 
to  stand  until  absorption  is  ended,  compare  this 
product  with  that  of  last  experiment. 

Ex.  41.  EtJiylene  (Oleflant  Gas). — Pour  into  a 
fifty-cubic-centimetre  flask  five  cubic  centimetres 
of  high-proof  alcohol,  and  then  add  slowly  twenty 
cubic  centimetres  of  strong  sulphuric  acid.  Con- 
nect with  a  pneumatic  trough  and  heat  carefully, 
protecting  the  glass  by  interposing  asbestos  paper 
between  the  flask  and  the  lamp.  Ethylene,  the 
gas  thus  obtained,  burns  with  a  brilliant  flame 


OLEFIANT  GAS.  79 

and  is  one  of  the  constituents  of  illuminating  gas. 
The  sole  products  of  its  combustion  are  carbonic 
dioxide  and  water,  and  it  must  therefore  contain 
both  carbon  and  hydrogen.  It  is  here  selected  as 
an  example  of  a  very  large  class  of  substances 
called  hydrocarbons.  The  phenomena  attending 
their  combustion  have  already  been  discussed 
(Ex.  28).  In  illuminating  gas,  a  very  complex 
product,  ethylene  is  mixed,  among  other  things, 
with  a  very  large  amount  of  another  hydrocarbon, 
called  me  than  (or  marsh  gas),  which  contains  only 
half  as  much  carbon  and  has  far  less  illuminating 
power.  The  petroleums  and  the  products  ob- 
tained from  them,  known  as  benzine,  kerosene, 
astral  oil,  paraffine,  etc.,  are  chiefly  mixtures  of 
a  great  number  of  hydrocarbons  (gases,  liquids, 
and  solids),  resembling  in  their  chemical  relations 
marsh  gas,  and  classed  with  it  under  the  general 
designation  of  the  paraffines.  Olefiant  gas  pre- 
pared as  above  is  so  called  because  it  unites  di- 
rectly either  with  chlorine  or  bromine  to  form  a 
liquid  which  has  an  oily  aspect,  and  there  are 
several  other  hydrocarbons  formed  in  the  distilla- 
tion of  coal  which  resemble  olefiant  gas  in  this  re- 
spect and  are  classed  with  it  under  the  name  of 
the  olefines.  Then  there  is  a  hydrocarbon  con- 
taining only  one  half  as  much  hydrogen  as  olefi- 
ant gas,  which  is  formed  abundantly  when  a 
Bunsen  lamp  burns  at  the  base,  and  is  at  once 


80  LABORATORY  PRACTICE. 

recognized  by  its  unpleasant  odour.  This  hydro- 
carbon is  also  one  of  a  class  known  as  the  acety- 
lenes, and  is  itself  usually  called  by  the  same 
name.  Lastly,  there  is  a  very  remarkable  class  of 
hydrocarbons  obtained  by  the  distillation  of  coal 
tar,  of  which  benzol  and  toluol  are  the  chief  mem- 
bers and  from  which  the  aniline  dyes  are  pro- 
duced. These  three  classes,  although  the  most 
important  groups,  by  no  means  include  all  the 
known  hydrocarbon  compounds,  while  the  possi- 
bilities of  multiplication  are  unlimited,  and  from 
the  great  family  of  hydrocarbons  the  almost  end- 
less products  of  organic  chemistry  may  be  de- 
rived. The  points  here  suggested  the  teacher  will 
expand  as  he  sees  fit. 

(1)  Charcoal  graphite  or  diamond  when  burnt  in  oxy- 
gen gas  all  yield  the  same  product  (carbonic  dioxide).     Are 
they  the  same  substance  ? 

(2)  Does  a  solution  of  carbonic  acid  in  water  conform  to 
your  general  conception  of  an  acid  ? 

(3)  Compare  the  composition  of  carbonic  oxide  and  car- 
bonic dioxide  with,  that  of  sulphurous  oxide  and  sulphuric 
oxide.     Do  all  these   oxides   form  acids  by  uniting  with 
water  ?    Regarding  the  intensity  of  the  acid  taste  and  of  the 
acid  reaction  as  some  measure  of  the  strength  of  the  acids^ 
what  would  be  your  estimate  of  the  relative  strength  of  the 
acids  thus  formed,  and  how  does  this  degree  of  strength  com- 
pare with  the  apparent  attraction  of  the  oxides  for  water  ? 

(4)  Is  the  proof  here  given  that  olefiant  gas  is  com- 
posed of  carbon  and  hydrogen  satisfactory  ?    Is  it  equally 
clear  that  the  gas  consists  only  of  carbon  and  hydrogen  ? 


PREPARATION   OF  NITRIC  ACID.  81 

8.  Nitrogen. 

Ex.  42.  Preparation  of  Nitrogen  Gas. — The 
aeriform  product  left  in  the  jar  from  Ex.  24  (b)  is 
nitrogen  gas.  The  student  should  test  the  gas  by 
immersing  in  it  a  lighted  match  ;  but  the  element- 
ary student  can  not  be  expected  to  learn  through 
actual  experiments  the  complex  relations  of  this 
remarkable  substance.  It  should,  however,  be 
made  evident  by  the  teacher  that  the  inability  to 
support  combustion  is  in  entire  harmony  with  the 
general  inert  relations  of  nitrogen  gas,  and  that 
this  is  its  most  striking  characteristics.  Never- 
theless, when  the  necessary  conditions  are  fulfilled, 
nitrogen  readily  enters  into  combination,  espe- 
cially with  oxygen  and  hydrogen,  forming  a  nu- 
merous and  important  class  of  products.  In  illus- 
tration of  this  last  point  the  formation  of  nitre 
should  be  explained  ;  and,  starting  from  this  nat- 
ural product,  the  student  may  prepare  a  few  well- 
marked  nitrogen  compounds  as  follows  : 

Ex.  43.  (a)  Preparation  of  Nitric  Acid. — Mix 
in  the  body  of  a  small  glass  retort  30  grammes  of 
nitre  with  the  same  weight  of  sulphuric  acid. 
Allow  the  mixture  to  stand  for  several  hours,  and 
then  distil  over  15  grammes  of  a  yellowish  liquid, 
which  is  nitric  acid  (a  very  important  chemical 
agent,  consisting  of  nitrogen  combined  with  both 
oxygen  and  hydrogen).  The  yellow  colour  of  the 


82  LABORATORY  PRACTICE. 

product  is  due  to  an  admixture  of  another  nitro- 
genized  product,  which  is  very  volatile,  and  may 
be  driven  off  by  gently  heating  the  acid  in  a 
flask. 

(b)  Nitric  Acid  contains  Oxygen.  — Take  in  a 
test  tube  four  or  five  cubic  centimetres  of  the  acid 
just  made.  Drop  into  it  in  small  portions  at  a 
time  not  over  one  decigramme  of  coarsely  pulver- 
ized roll  brimstone.*  Cautiously  heat  to  boiling 
and  maintain  gentle  ebullition  for  some  time. 
Largely  dilute  with  water  and  decant  into  a  clean 
test  tube  from  the  remaining  sulphur.  Test  with 
solution  of  baric  chloride  as  described  in  Ex.  33. 
It  will  thus  appear  that  sulphuric  acid  is  formed  by 
action  of  nitric  acid  on  sulphur.  Remembering 
now  that  sulphuric  acid  consists  of  sulphur,  oxy- 
gen, and  water,  draw  your  own  inferences. 

Indeed,  nitric  acid  contains  so  much  oxygen, 
held  feebly  in  combination,  that  it  furnishes  an 
efficient  means  of  uniting  oxygen  to  other  bodies. 
It  may  sustain  combustion  like  the  atmosphere, 
and  it  is  for  these  reasons  said  to  be  a  powerful 
oxidizing  agent.  The  teacher  can  effectively  illus- 
trate this  point  by  pouring  from  a  long-handled 
glass  or  porcelain  spoon  a  few  cubic  centimetres 
of  the  strongest  acid  on  finely  pulverized  charcoal. 
The  powder  should  be  well  dried  by  heating  it  to 

*  The  action  of  very  strong  nitric  acid  on  combustible  matter 
is  often  violent,  and  this  experiment  should  be  made  with  caution. 


COMPOSITION  OF  NITRIC  ACID.  83 

incipient  redness  in  a  porcelain  dish,  and  while 
still  hot  the  acid  poured  upon  it  (a  few  drops  at 
a  time).  The  charcoal  will  then  flash  almost  like 
gunpowder.  In  this  connection  the  teacher  may 
add  that  in  the  explosion  of  gunpowder  charcoal 
burns  at  the  expense  of  the  oxygen  stored  in  the 
nitre. 

(c)  Nitric  Acid  contains  Nitrogen  and  Water. 
—This  can  be  shown  by  passing  the  vapor  of  the 
acid  carried  by  a  current  of  carbonic  dioxide  over 
copper  clippings  heated  to  redness  in  a  combustion 
tube.  The  metal  will  take  up  all  the  oxygen  in 
the  acid  except  that  belonging  to  the  water  present, 
while  the  water  set  free  may  be  collected  by  pass- 
ing the  current  after  leaving  the  combustion  tube 
through  a  U-tube  packed  in  ice.  If,  lastly,  the 
current  is  passed  on  to  a  small  pneumatic  trough 
filled  with  water  holding  caustic  alkali  in  solu- 
tion the  carbonic  dioxide  will  be  absorbed  and 
nitrogen  gas  collected.  A  regulated  current  of 
carbonic  dioxide  is  easily  obtained  with  the  gen- 
erator described  in  note  to  Ex.  25,  tilling  the  bell 
with  broken  marble  and  the  beaker  with  common 
muriatic  acid  diluted  with  an  equal  volume  of 
water.  The  few  cubic  centimetres  of  nitric  acid 
required  are  best  held  in  a  bulb  tube  so  support- 
ed that  it  may  be  warmed  with  a  lamp,  and  con- 
nected at  one  end  with  the  generator  and  at  the 
other  with  the  combustion  tube  by  a  rubber  con- 


84  LABORATORY   PRACTICE. 

nector  (corks  would  be  instantly  corroded).  The 
combustion  tube  may  be  arranged  as  in  Ex.  33, 
but  should  be  somewhat  larger,  and  is  best  filled 
with  finely  pulverized  copper,  such  as  is  obtained 
by  the  reduction  of  copper  oxide.  The  small 
pneumatic  trough  is  easily  extemporized  out  of  a 
glass  or  porcelain  dish  and  a  large  test  tube. 
Other  details  of  the  apparatus  may  now  be  left  to 
the  ingenuity  of  the  student ;  but  the  experiment 
should  not  be  intrusted  except  to  skilful  manipu- 
lators, and  in  most  cases  will  best  be  shown  on  the 
lecture  table.  In  all  cases,  however,  he  should 
make  a  sketch  of  the  apparatus  in  his  note  book 
and  point  out  the  use  of  each  part  and  justify  the 
conclusion  that 

Nitrogen,  oxygen,  and  water  yield  nitric  acid. 

(1)  Is  the  proof  which  has  been  given  of  the  composi- 
tion of  nitric  acid  synthetical  or  analytical  ?    The  synthesis 
of  nitric  acid  can  not  be  readily  made  because,  in  conformity 
to  its  great  inertness,  nitrogen  does  not  combine  directly 
with  oxygen.     Nevertheless,  by  indirect  means,  the  oxide  of 
nitrogen  corresponding  to  nitric  acid  has  been  prepared.     It 
is  a  white  crystalline  solid  which  eagerly  unites  with  water. 
Compare  sulphuric  oxide. 

(2)  You  have  now  handled  the  most  important  acids 
used  in  a  chemical  laboratory.     Make  a  list  of  them  with 
the  composition  of  each  so  far  as  you  have  discovered  it. 
Make  clear  that  you  can  recognize  all  of  them  whether  con- 
centrated or  diluted.     In  every  case  a  specific  test  has  not 
been  given,  but  by  inquiry  of  your  teacher  or  elsewhere  seek 
the  necessary  knowledge  until  you  are  perfectly  certain  of 
your  ability  to  distinguish  all  these  substances. 


NITRIC  OXIDE.  85 

Ex,  44.  (a)  Preparation  of  Nitric  Oxide.— 
Place  fifty  grammes  of  copper  clippings  in  a  glass 
flask  fitted  with  a  cork,  through  which  passes  a 
tube  funnel  as  well  as  an  evolution  tube.  Drench 
the  clippings  with  water,  and  then  pour  on 
through  the  funnel  in  successive  portions  the 
product  of  the  last  experiment  mixed  with  three 
times  its  volume  of  water.  After  the  action  has 
started  add  a  fresh  portion  from  time  to  time  as 
the  effervescence  slackens.  Connect  with  a  pneu- 
matic trough  and  collect  the  gas  in  a  quick-sealing 
jar.  The  colorless  gas  is  a  compound  of  nitrogen 
and  oxygen,  called  nitric  oxide.  The  deep -red 
fumes  which  appear  in  the  generating  flask,  and 
whenever  the  nitric  oxide  mixes  with  the  oxygen 
gas  of  the  air,  is  another  compound  of  nitrogen 
and  oxygen  containing  more  oxygen  and  called 
nitric  peroxide  ;  and  it  is  chiefly  this  adventitious 
product  which  imparts  the  yellow  tint  to  the 
crude  nitric  acid. 

(b)  Analysis  of  Nitric  Oxide. — In  a  pint  jar 
of  nitric  oxide  collected  in  the  last  experiment 
burn  a  small  bit  of  phosphorus,  not  exceeding  one 
fourth  of  a  gramme  in  weight,  with  all  the  pre- 
cautions stated  in  Ex.  24.  Notice  that  the  same 
white  product  is  formed  as  when  phosphorus 
burns  in  pure  oxygen  gas  or  in  air.  Hence  we 
may  infer  that  nitric  oxide  contains  oxygen. 
After  the  jar  is  cold  open  the  mouth  under  water. 


86  LABORATORY  PRACTICE. 

Notice  that  the  residual  gas  fills  only  one  half  of 
the  volume  of  the  jar,  and,  further,  that  it  has  the 
characteristic  inertness  of  nitrogen.  Hence  the 
additional  conclusion  that  nitric  oxide  consists  of 
equal  volumes  of  nitrogen  gas  and  oxygen  gas. 

(1)  Compare  the  action  of  copper  on  nitric  acid  in  the 
last  two  experiments.  Notice  that  while  in  the  first  it  re- 
duces the  acid  to  water  and  nitrogen  gas,  in  the  second  it 
does  not  remove  so  much  oxygen  and  leaves  nitric  oxide. 
Observe  also  that  nitric  oxide  shows  no  tendency  to  unite 
with  water. 

Ex.  45.  Preparation  of  Ammonia. — Mix  in  a 
gasometer  *  one  volume  of  nitric  oxide  with  two 
and  a  half  volumes  of  hydrogen,  and  pass  a  slow 

*  A  very  useful  and  inexpensive  gasometer  may  be  made  of  a 
large  glass  bottle  of  the  capacity  of  two  quarts  or  one  gallon,  as  re- 
quired. Fit  tightly  in  the  neck  a  rubber  cork  with  three  perfora- 
tions. Through  these  perforations  pass  glass  tubes,  all  bent  at  right 
angles  a  short  distance  above  the  cork.  Two  of  the  tubes  should 
reach  the  bottom  of  the  bottle,  the  third  only  pass  through  the  stop- 
per. Connect  one  of  the  longer  tubes  with  a  sink  by  a  length  of 
rubber  hose.  In  the  same  way  connect  the  second  of  the  longer 
tubes  with  a  water  tap ;  and,  lastly,  slip  on  to  the  shorter  tube  a 
third  length  of  rubber  hose  to  serve  as  an  outlet  for  the  gas.  Guard 
all  three  of  the  tubes  with  pressure  taps.  Begin  by  filling  the  ga- 
someter with  water,  allowing  the  liquid  to  flow  in  from  the  tap  and 
the  air  to  escape  from  the  outlet.  When  full,  close  the  outlet,  open 
the  overflow  into  the  sink,  remove  the  rubber  hose  from  the  water 
tap,  and  connect  it  with  the  gas  generator.  The  gas  will  then  flow 
in,  the  displaced  water  running  off  by  the  overflow.  When  a  suffi- 
cient amount  of  gas  has  been  collected  the  overflow  must  be  closed 
and  the  rubber  hose  removed  from  the  generator  and  replaced  on 
the  water  tap.  Then,  on  opening  the  tap,  water  will  flow  in  again 
and  drive  out  the  gas  by  the  outlet  tube  through  any  apparatus 
with  which  it  may  be  connected. 


AMMONIA.  87" 

stream  of  this  mixture  through  a  short  tube  filled 
with  platinized  asbestos  and  heated  to  a  low  red 
heat.  There  will  be  an  abundant  formation  of 
aqueous  vapours,  indicating  that  a  combination 
has  taken  place  between  the  hydrogen  gas  and 
the  oxygen  of  the  nitric  oxide ;  and  at  the  same 
time  there  will  be  developed  a  strong  pungent 
odour,  familiar  to  every  one  as  the  odour  of  am- 
monia, which  must  evidently  be  formed  by  the 
union  of  nitrogen  with  hydrogen.  This  pungent 
product  is  a  gas,  and  common  aqua  ammonia  is  a 
solution  of  the  gas  in  water.  To  obtain  a  more 
familiar  acquaintance  with  this  important  chem- 
ical agent,  half  fill  a  small  glass  flask  with  con- 
centrated aqua  ammonia  and  close  the  neck  with 
a  rubber  stopper,  through  which  passes  an  evolu- 
tion tube  leading  to  the  top  of  a  quick- sealing 
glass  jar,  arranged  as  in  Ex.  26.  Heat  the  liquid 
in  the  flask  to  boiling,  when  a  large  volume  of 
colourless  gas  will  pass  over  and  may  be  collected 
by  displacement.  When  the  jar  is  full,  the  over- 
flow will  be  at  once  recognized  by  the  strong 
odour.  The  process  may  then  be  stopped,  the  jar 
closed,  and  the  gas  preserved  for  another  experi- 
ment. 

Ex.  46.  (a)  Ammonia  Salts. — Mix  five  cubic 
centimetres  of  strong  aqua  ammonia  with  twice  its 
volume  of  water.  In  the  same  way  mix  five  cubic 
centimetres  of  strong  nitric  acid  with  twice  its 


88  LABORATORY  PRACTICE. 

volume  of  water.  Study  the  effects  of  these  solu- 
tions on  test  papers,  both  litmus  and  turmeric. 
Next  add  slowly  the  ammonia  to  the  nitric  acid 
until  the  opposite  effects  exactly  neutralize  each 
other.  Lastly,  evaporate  the  mixture  at  a  low 
heat  until  a  drop  taken  out  on  a  rod  solidifies, 
and  then  when  the  dish  is  set  on  one  side  a  white 
salt,  called  ammonic  nitrate,  will  crystallize  out. 

(b)  Make  the  same  experiment  with  sulphuric 
acid,  with  phosphoric  acid,  with  hydrochloric 
acid,  and  with  carbonic  acid.  In  the  last  case 
add  the  fifteen  cubic  centimetres  of  diluted  aqua 
ammonia  to  a  jar  of  carbonic-acid  gas  and  shake 
together.  Evaporate  each  solution  to  dryness  and 
collect  the  crystalline  salt. 

Obviously  the  solution  of  ammonia  gas  (aqua 
ammonia)  sustains  relations  which  are  the  very 
opposite  to  those  of  the  acids,  and  belongs  to  an 
equally  important  class  of  chemical  products 
variously  called  alkalis  (when  the  solutions  are 
caustic)  or  bases  (when  they  are  not).  The  acids 
exhibit  no  tendency  to  unite  with  each  other,  but 
they  eagerly  unite  with  the  anti-acids  (as  in  the 
above  experiments)  to  form  a  class  of  bodies,  for 
the  most  part  crystalline,  called  salts.  The  anti- 
acids  are  formed,  as  a  rule,  by  the  union  of  met- 
als with  oxygen  and  water,  while  the  acids  (as 
we  have  seen)  result  as  generally  from  a  similar 
union  of  bodies,  like  phosphorus,  sulphur,  and 


MAGNESIUM  AND  ZINC.  89 

carbon,  which  do  not  exhibit  metallic  characters. 
Having  studied  the  relations  of  a  few  of  these  so- 
called  metalloids  we  pass  next  to  study  the  rela- 
tions of  several  metals. 


9.  Magnesium  and  Zinc. 

Ex.  47.  Specific  Characters.  —  The  student 
should  be  given  a  gramme  of  magnesium  ribbon, 
and  with  this  he  should  study  the  characters  of 
the  metal  (colour,  lustre,  tenacity,  specific  gravity) 
and  compare  these  with  the  corresponding  prop- 
erties of  metallic  zinc,  also  given  to  him  rolled  out 
into  ribbon  of  about  the  same  size.  Reserving  a 
short  length  of  the  magnesium  ribbon  for  burn- 
ing, the  student  should  dissolve  the  rest  in  a  few 
cubic  centimetres  of  dilute  sulphuric  acid.  Col- 
lect and  examine  the  gas  evolved,  and  compare 
the  reaction  with  that  in  Ex.  25.  Lastly,  let  him 
evaporate  the  solution  of  magnesium  thus  ob- 
tained on  a  watch  glass,  and  compare  the  crystal- 
line residue  with  that  obtained  from  the  solution 
of  zinc  formed  in  the  above- cited  experiment. 

Ex.  48.  Burning  of  Magnesium.  —  Let  the 
student  burn  the  reserved  piece  of  magnesium 
ribbon,  holding  it  by  pincers  and  lighting  it  like 
a  match,  and  let  him  compare  the  combustibility 
and  colour  of  the  flame  with  that  produced  with  a 
similar  ribbon  of  zinc.  In  order  that  he  may  fur- 


90  LABORATORY   PRACTICE. 

ther  study  the  nature  of  the  product  formed,  fur- 
nish the  student  with  half  a  gramme  of  magne- 
sium* powder.  Let  him  place  this  on  a  small 
square  of  asbestos  paper  (previously  ignited)  and 
weigh  the  amount  accurately  on  the  pan  of  a 
balance.  Let  him  now  ignite  the  powder  with  a 
match,  and  when  burned  out  and  cooled  reweigh 
it.  What  means  the  increase  of  weight,  and  what 
must  be  the  composition  of  the  white  powder 
left?  Transfer  the  powder  to  a  small  evaporat- 
ing dish.  Thoroughly  drench  with  water.  Place 
a  small  bit  of  the  wet  powder  on  red  litmus 
paper.  Compare  the  eifect  with  that  of  an  acid. 
Lastly,  dissolve  the  residue  in  the  smallest  pos- 
sible amount  of  dilute  sulphuric  acid,  adding 
the  acid  drop  by  drop.  Evaporate  the  solu- 
tion till  a  crust  appears  and  leave  to  crystallize. 
Can  you  recognize  the  saline  product  by  the 
taste  ? 

Try  the  same  experiments  with  zinc ;  but  its 
powder  does  not  burn  so  readily,  and  it  is  more 
difficult  to  recognize  the  increase  of  weight.  Nev- 
ertheless, it  is  easily  burned  by  sifting  the  powder 
on  to  a  sheet  of  paper  through  the  flame  of  a 
Bunsen  burner  held  obliquely,  or  by  spreading 

*  A  few  years  ago  magnesium  would  have  been  too  expensive 
for  general  experimenting,  but,  as  a  result  of  its  application  in  the 
arts,  it  can  now  be  purchased  for  a  moderate  price.  At  the  price 
quoted  by  German  dealers  in  chemicals  the  two  grammes  required 
for  each  student  would  cost  less  than  three  cents. 


MAGNESIUM  AND  ZINC.  91 

the  powder  over  a  square  of  asbestos  paper  and 
playing  on  it  with  the  same  flame. 

(1)  The  comparison  of  two  metals  closely  resembling 
each  other,  like  magnesium  and  zinc,  affords  excellent  prac- 
tice, and  may  be  used  to  test  the  student's  skill  in  observa- 
tion and  deduction.    If  further  practice  is  thought  necessary 
a  comparison  may  be  made  between  two  metals  resembling 
each  other  still  more  closely;   as,  for  example,  iron  and 
nickel.     In  all  such  cases  the  student  should  be  required  to 
work  out  the  results  unaided,  and  make  a  full  and  clear 
statement  in  his  note  book  of  what  he  observes  and  what  he 
infers.     His  work  should  then  be  carefully  criticised,  and,  if 
necessary,  the  experiments  repeated,  after  fresh  directions  or 
suggestions  from  the  teacher.     The  student  should  be  led  to 
appreciate  the  fact  that  although  the  distinctions  between 
substances  are  usually  broad  and  clear,  they  are  also  at  times 
narrow  and  indefinite,  and  that  the  identification  or  differ- 
entiation of  a  newly  found  substance  often  turns  on  minute 
observations  and  delicate  discriminations. 

(2)  How  can  you  prove  that  magnesic  oxide  combines 
with  water  ?    Does  zinc  oxide  combine  in  like  manner  ? 

Compare — 

Magnesium,  oxygen,  and  water  yield  magnesic  hydrate 
—basic. 

Sulphur,  oxygen,  and  water  yield  sulphuric  acid — acid. 

Further,  it  appears  that — 

Magnesic  hydrate  and  sulphuric  acid  yield  Epsom  salts 
—salt. 

Would  magnesic  oxide  yield  the  same  product  as  mag- 
nesic hydrate  ? 

(3)  Compare  the  product  of  the  action  of  magnesium  on 
dilute  sulphuric  acid  with  that  obtained  by  dissolving  mag- 
nesic oxide  in  the  same  solvent.     Why  is  hydrogen  gas  not 
evolved  in  the  second  process  ? 


92  LABORATORY   PRACTICE. 

10.  Sodium. 

Ex.  49.  Specific  Characters. — The  specific  char- 
acters of  this  interesting  alkaline  metal  should  be 
shown  as  far  as  possible  to  the  student ;  but  it 
will  seldom  be  advisable  to  intrust  the  material  to 
inexperienced  hands,  and  equally  good  practice 
in  studying  specific  characters  can  be  had  with 
cheaper  and  less  dangerous  substances.  The  ac- 
tion of  sodium  on  water  illustrates  principles  so 
fundamental  in  the  theory  of  chemistry  that  the 
experiment  should  on  no  account  be  omitted. 
The  action  of  the  pure  metal  is  always  violent, 
and  frequently  dangerous  ;  still,  it  is  a  very  inter- 
esting experiment,  which  the  teacher  may  make 
before  the  class,  with  proper  precautions.  The 
best  way  is  to  throw  a  bit  of  sodium,  not  larger 
than  a  pea,  on  some  sheets  of  porous  paper  thor- 
oughly soaked  and  running  with  water.  The 
melted  globule  is  thus  prevented  from  swimming 
round,  and  the  heat  developed  by  the  chemical 
change  accumulates  to  such  an  extent  as  to  in- 
flame the  escaping  hydrogen,  which  burns  with  a 
flame  that  is  coloured  yellow  by  the  presence  of 
sodium  vapour.  For  use  of  students,  an  amalgam 
of  sodium — one  part  of  sodium  to  about  ten  parts 
of  mercury — should  be  prepared  ;  and  this  may  be 
used  with  entire  safety.  The  action  of  sodium  on 
mercury  is  violent,  but  the  amalgam  can  be  easily 


ALKALIES.  93 

made  by  heating  the  mercury  to  about  200°  in  a 
Hessian  crucible  of  eight  or  ten  times  the  capacity 
required  to  hold  the  metal,  and  then  adding  the 
sodium  in  one  large  bar.  Assuming  this  amal- 
gam to  have  been  previously  prepared  or  pur- 
chased, the  student  should  make  the  following 
experiment,  having,  of  course,  been  previously  told 
the  object  of  using  the  mercury,  and  that  it  plays 
no  part  in  the  chemical  change  :  Take  a  small  flask 
fitted  for  the  evolution  of  gas,  and  place  in  it 
about  twenty-five  grammes  of  water  ;  add  now  a  few 
lumps  of  the  amalgam,  and  collect  the  gas  over 
the  pneumatic  trough.  As  the  evolution  lessens, 
hasten  it  by  heating  the  flask  with  a  lamp.  Burn 
the  gas  that  is  collected,  and  recognize  that  it  is 
hydrogen.  Pour  off  now  the  solution  left  in  the 
flask  from  the  mercury,  and,  in  the  first  place, 
test  it  with  litmus  and  turmeric  paper  which 
have  previously  been  dipped  in  a  very  weak  acid. 
It  will  thus  be  seen  that  the  product  is  a  sub- 
stance which,  like  the  solution  of  ammonia,  re- 
verses the  effect  of  an  acid  on  vegetable  dyes ;  in 
other  words,  that  it  is  basic.  Such  soluble  and 
caustic  bases  are  called  alkalies.  Rub  a  few  drops 
of  the  liquid  between  the  fingers,  and  notice  the 
effect,  which  is  termed  caustic.  Evaporate  now 
the  liquid,  and  compare  the  residue  with  caustic 
soda.  Redissolve  this  residue  in  a  very  small 
amount  of  water,  and  divide  the  solution  between 
7 


94:  LABORATORY  PRACTICE. 

three  watch,  glasses  ;  neutralize  the  solution  in  the 
first  glass  with  a  few  drops  of  hydrochloric  acid, 
that  in  the  second  glass  with  a  few  drops  of  nitric 
acid,  and  add  the  contents  of  the  third  watch  glass 
to  about  twenty -five  cubic  centimetres  of  carbonic 
acid  (soda  water).  Allow  the  solutions  to  evapo- 
rate, and  examine  the  crystals  formed  with  a  lens; 
also  attempt  to  recognize  the  products  by  tasting 
the  residue  in  each  case  ;  they  will  be  discovered 
to  be  common  salt,  sodic  nitrate,  and  sodic  car- 
bonate, respectively.  As  caustic  soda  has  thus 
been  made  solely  with  sodium  and  water,  a  prob- 
able inference  in  regard  to  its  composition  may  at 
once  be  drawn,  and  the  three  familiar  products 
last  obtained  will  be  recognized  as  salts  of  sodi- 
um. In  this  connection  the  student  should  be 
told  about  the  sources  of  these  substances,  and 
their  uses  in  daily  life  and  in  the  arts. 

Ex.  50.  (a)  Using  a  small  amount  of  magnesi- 
um or  zinc  powder  and  boiling  with  water  in  a 
test  tube,  compare  the  action  of  these  metals  on 
water  with  that  of  sodium.  It  will  appear  that — 

Sodium  acts  violently  on  water  and  yields  so- 
dic hydrate  and  hydrogen  gas. 

Magnesium  acts  slowly  on  water  and  yields 
magnesic  hydrate  and  hydrogen  gas. 

Zinc  acts  very  feebly  on  water  and  yields  zinc 
oxide  and  hydrogen  gas. 

(b)  The  teacher  should  burn  sodium  in  dry 


METALLIC  OXIDES.  95 

oxygen,  first  melting  the  metal  in  an  iron  spoon. 
Dissolve  the  oxide  formed  in  water,  evaporate  to 
dryness,  and  compare  it  with  product  of  the  direct 
action  of  sodium  on  water.  Care  must  be  taken 
to  separate  the  white  powder  from  unburnt  metal 
before  adding  water,  and  the  experiment  should 
not  be  intrusted  to  unskilful  students.  It  will 
now  further  appear  that— 

Sodium  unites  with  oxygen  to  form  sodic  oxide 
(white  powder). 

Magnesium  unites  with  oxygen  to  form  mag- 
nesic  oxide  (white  powder). 

Zinc  unites  with  oxygen  to  form  zinc  oxide 
(white  powder). 

Also  that — 

Sodium  oxide  unites  with  water  eagerly  to 
form  sodic  hydrate. 

Magnesium  oxide  unites  with  water  feebly  to 
form  magnesic  hydrate. 

Zinc  oxide  will  not  unite  with  water. 

The  striking  differences  depend  on  the  fact 
that  sodic  hydrate  is  very  soluble  in  water.  In- 
terpret all  the  phenomena. 

(c)  The  teacher  should  burn  sodium  in  dry 
chlorine.  It  burns  readily  when  heated  above 
melting  point  in  an  iron  spoon,  but  the  experi- 
ment should  be  made  under  a  hood  with  powerful 
draught.  A  considerable  part  of  the  product  vola- 
tilizes. This  is  readily  dissolved  in  water  and  crys- 


96  LABORATORY  PRACTICE. 

tallized,  when  the  cubic  form  of  the  crystal  and 
taste  show  it  to  be  common  salt.  It  thus  appears 
that  sodium  and  chlorine  yield  common  salt.  So- 
dic  hydrate  (or  sodic  oxide)  and  hydrochloric  acid 
also  yield  common  salt.  Consider  in  what  respect 
the  formation  of  common  salt  differs  from  that  of 
other  salts — for  example,  magnesium  sulphate. 

Ex.  51.  Volatile  and  Fixed  Alkali. — Compare 
the  formation  of  the  ammonia  with  that  of  the 
sodium  salts. 

Ammonia,  gas,  water,  and  sulphuric  acid  yield 
ammonic  sulphate. 

Ammonia,  gas,  water,  and  nitric  acid  yield 
ammonic  nitrate. 

Ammonia,  gas,  water,  and  hydrochloric  acid 
yield  ammonic  chloride. 

Sodic  hydrate  and  sulphuric  acid  yield  sodic 
sulphate. 

Sodic  hydrate  and  nitric  acid  yield  sodic  ni- 
trate. 

Sodic  hydrate  and  hydrochloric  acid  yield 
sodic  chloride. 

Compare  the  two  salts  of  each  acid  and  deter- 
mine whether  volatile  or  not.  Heat  solutions  of 
each  of  the  ammonia  salts  with  a  solution  of  caus- 
tic soda  and  test  the  gas  given  off  with  red  litmus 
paper  and  with  great  caution  by  the  smell. 


REDUCTION  AND  OXIDATION.  97 

11.  Copper. 

Ex.  52.  Distinctive  Characters. — The  metal  is 
best  used  in  a  very  thin,  flexible  sheet,  easily  cut 
with  a  pair  of  scissors.  Let  the  student  first  study 
and  describe  the  properties  of  the  metal,  deter- 
mining its  specific  gravity  in  the  usual  way,  and 
comparing  its  colour,  hardness,  toughness,  mallea- 
bility, etc.,  with  the  similar  qualities  of  zinc.  Let 
him  harden  on  an  anvil  and  anneal  with  heat.  Let 
him  next  heat  in  separate  test  tubes  a  few  bits  of 
the  metal  with  nitric,  hydrochloric,  and  sulphuric 
acids,  using  both  strong  and  weak  acid.  Com- 
pare with  zinc.  Observe  behaviour,  to  be  inter- 
preted beyond. 

Ex.  53.  (a)  Reduce  Oxide  with  Hydrogen  Gas 
and  reoxidize  in  the  Air. — Introduce  into  a  com- 
bustion tube  about  twenty  grammes  of  black  ox- 
ide of  copper.  Take  the  tare  on  the  balance. 
Mount  on  tube  furnace,  connecting  one  end  with 
a  hydrogen  generator  and  the  other  end  with  a 
small  U  tube  kept  cool  with  ice.  Pass  now  a  slow 
current  of  hydrogen  gas  over  the  powder,  and 
after  the  gas  has  expelled  the  air  heat  the  com- 
bustion tube  to  low  redness.  Observe  that  the 
powder  takes  on  the  colour  and  lustre  of  metallic 
copper  and  that  water  collects  in  the  tube.  After 
the  reduction  is  complete,  dismount  the  apparatus 
and  reweigh  the  tube. 


98  LABORATORY  PRACTICE. 

(b)  Kemount  on  the  furnace  the  combustion 
tube  with  its  contents  as  left  in  the  last  experi- 
ment. Leave  one  end  open  to  the  air  and  con- 
nect the  other  with  the  gasometer  before  de- 
scribed (note  to  Ex.  45),  which,  having  been  pre- 
viously filled  with  water  and  the  overflow  opened 
into  a  sink  at  a  lower  level,  will  act  as  an  aspira- 
tor. Regulate  by  pressure  tap  so  that  air  will  be 
drawn  through  the  tube  not  faster  than  about  two 
bubbles  a  second,  and  then  heat  the  reduced  cop- 
per to  redness.  After  collecting  one  or  two  litres 
of  gas  close  and  dismount  the  gasometer,  while  at 
the  same  time  connecting  the  combustion  tube 
with  an  aspirator  pump  to  hasten  the  process. 
Continue  heating  the  tube  until  the  powder  has 
acquired  a  uniform  black  color.  Then  allow  to 
cool,  dismount,  and  reweigh.  Meanwhile  transfer 
the  gas  collected  to  gas  bottles  over  a  pneumatic 
trough  and  seek  to  recognize  the  substance.  In- 
terpret all  the  phenomena  observed.  Compare 
Ex.  24  (b)  and  Ex.  42.  Keep  the  oxide  of  copper 
in  the  tube  for  another  experiment.  . 

(1)  Water  consists  of  hydrogen  and  oxygen.     The  pro- 
duction of  water  from  hydrogen  gas  implies  what?    The 
formation  of  water  is  attended  with  the  change  of  the  black 
powder  to  metallic  copper.     Of  what  must  that  powder  con- 
sist? 

(2)  Whence  comes  the  nitrogen  collected  in  the  gasome- 
ter ?    How  may  free  oxygen,  or  substances  holding  oxygen 
loosely  united,  be  expected  to  act  on  free  copper  ?    How  may 


SALTS  OF  COPPER.  99 

oxide  of  copper  be  expected  to  act  on  substances  containing 
hydrogen  or  carbon  ? 

Ex.  54.  (a)  Salts  of  Copper. — Take  three  test 
tubes  holding  a  few  cubic  centimetres  of  dilute 
sulphuric,  hydrochloric,  and  nitric  acid  respect- 
ively. Dissolve  in  each  black  oxide  of  copper, 
adding  in  very  small  quantities  so  long  as  solution 
is  obtained  on  boiling.  Evaporate  on  watch  glasses 
and  crystallize  the  products. 

Oxide  of  copper  and  sulphuric  acid  give  cop- 
per sulphate  (salt). 

Oxide  of  copper  and  nitric  acid  give  copper 
nitrate  (salt). 

Oxide  of  copper  and  hydrochloric  acid  give 
copper  chloride  (salt). 

(&)  Take  a  few  cubic  centimetres  of  the  solution 
of  copper  in  nitric  acid — Ex.  52  or  44  (a).  Secure 
the  ready  solution  of  copper  in  sulphuric  and  hy- 
drochloric acids  by  adding  a  few  drops  of  nitric 
acid  in  each  case  ;  evaporate  to  dryness  (not  over 
100°),  dissolve  the  residue  in  the  smallest  possible 
amount  of  water,  and  allow  to  crystallize  in  a 
warm  place.  Compare  products  with  those  ob- 
tained in  (a).  Lastly,  heat  copper  nitrate  in  a 
porcelain  crucible  to  low  redness,  until  fumes 
cease,  then  to  bright  redness.  When  cold  pulver- 
ize the  residue  in  a  mortar  and  examine  the  black 
powder  left.  Seek  to  interpret  now  the  phenomena 
observed  when  attempting  to  dissolve  the  metals. 


100  LABORATORY  PRACTICE. 

(i)  What  is  the  difference  between  the  action  of  copper 
and  zinc  on  dilute  sulphuric  acid  ?  This  difference  can 
be  explained  by  the  thermal  relations  of  the  metals,  but 
must  here  be  accepted  as  a  fact.  Why  should  you  expect 
that  the  addition  of  nitric  acid  would  secure  the  ready  solu- 
tion of  copper  in  both  sulphuric  and  hydrochloric  acids  ? 
What  double  part  does  nitric  acid  play  in  dissolving  copper  ? 
Compare  Ex.  44  (a)  and  also  the  additional  fact  that  when 
in  this  preparation  of  nitric  oxide  the  temperature  is  allowed 
to  rise  too  high  and  the  action  to  become  violent  the  prod- 
uct is  chiefly,  or  even  altogether,  nitrogen  gas. 


12.  Iron. 

Ex.  55.  Distinguishing  Properties. — The  met- 
al is  best  used  in  the  shape  of  wrought-iron-wire 
nails  or  tacks  of  various  sizes,  also  in  fine  powder 
"iron  by  hydrogen."  Let  the  student  first  study 
and  describe  the  properties  of  the  metal,  as  in  the 
case  of  copper  (Ex.  52).  He  should  then  find  its 
specific  gravity,  using  a  method  applicable  in  many 
cases.  Take  a  small  vial  or  flask  (ten  to  twenty 
cubic  centimetres  capacity)  and  make  a  mark  on 
the  narrow  neck.  Fill  the  bottle  to  the  mark  with 
water  and  tare  it  on  the  balance.  Select  some 
iron  nails  as  large  as  will  conveniently  be  held  by 
the  bottle.  Place  about  twenty  grammes  of  these 
nails  at  the  side  of  the  bottle,  and  take  the  exact 
weight.  Remove  the  bottle  from  the  pan  and 
drop  in  the  nails  one  by  one,  taking  care  to  avoid 
entangling  bubbles  of  air.  Wipe  off  the  water 
which  runs  over,  and  with  a  small  roll  of  porous 


OXIDATION  OF  IHON.  101 

paper  reduce  the  level  in  the  neck  to  the  mark. 
Replace  on  the  pan  and  reweigh,  when  the 
loss  of  weight  is  obviously  the  weight  of  water 
displaced  by  the  nails.  In  this  connection  the 
qualities  and  relations  of  the  different  kinds  of 
iron — wrought  iron,  cast  iron,  and  steel — should 
be  explained  by  the  teacher,  and  the  prominent 
features  of  the  metallurgy  of  iron  might  appropri- 
ately be  discussed. 

Ex.  56.  (a)  Burning  of  Iron. — If  the  iron  pow- 
der is  sufficiently  fine  it  will  burn  on  the  pan  of  a 
balance  like  magnesium  (Ex.  48),  and  the  increase 
of  weight  may  be  found.  If  too  coarse  to  burn  in 
this  way,  the  iron  powder  will  burn  brilliantly  by 
sprinkling  it  through  the  flame  of  a  Bunsen  burn- 
er held  obliquely.  The  residue  may  be  collected 
on  a  large  sheet  of  paper  held  obliquely.  A  more 
striking  experiment  of  burning  a  watch  spring  in 
oxygen  gas  should  be  made  by  the  teacher.  After 
removing  the  temper  of  the  steel  by  heating  in  a 
lamp  flame,  the  spring  can  be  coiled  into  a  spiral. 
Tip  one  end  with  sulphur  like  a  match,  and  hang 
the  other  from  a  wire,  which  should  be  slid 
through  a  cork,  closing  the  tubulature  of  the  jar 
as  fast  as  the  oxygen  is  exhausted.  Protect  the 
exposed  face  of  the  cork  with  metallic  foil,  and 
then  light  the  iron  match  and  plunge  it  into  the 
gas.  The  experiment  is  best  made  in  a  tubulated 
bell  jar  standing  over  water  into  which  the  oxide 


102  LABORATORY  PRACTICE. 

of  iron  melted  by  the  flame  falls  in  drops.  These 
melted  globules  would  crack  the  bottom  of  a  glass 
jar  unless  protected  by  sand. 

(b)  Weigh  out  in  a    porcelain    crucible    five 
grammes  of  the  iron  powder.     Ignite  with  fre- 
quent stirring  (use  iron  wire)   until  completely 
oxidized.     Reweigh  when  cold,  and  interpret  the 
result. 

(c)  Weigh  out  five  grammes  of  iron  powder  in 
a  shallow  porcelain  dish.     Keep  the  powder  moist 
with  water  until  it  is  wholly  converted  into  rust. 
Then  allow  to  dry  in  the  air  and  weigh  again. 
Save  a  part  of  the  residue  for  further  use. 

(1)  Why  is  the  weight  of  the  rust  in  (c)  greater  than  that 
of  the  red  powder  formed  in  (6)  ?  Devise  an  experiment  for 
testing  your  inference. 

Ex.  57.  (a)  Iron  and  Sulphuric  Acid. — Iron 
dissolves  in  dilute  sulphuric  acid  like  zinc  or  mag- 
nesium, with  rapid  evolution  of  hydrogen  gas. 
Repeat  Ex.  25,  using  5  grammes  wrought-iron 
tacks  (instead  of  zinc  clippings),  also  10  grammes 
strong  sulphuric  acid  diluted  with  30  grammes 
of  water.  Use  a  100-cubic-centimetre  flask.  Col- 
lect the  gas  as  before,  and  compare  with  the  pre- 
vious product.  Observe  precautions  previously 
given  (note  to  Ex.  27)  in  regard  to  the  preparation 
of  hydrogen  gas. 

(b)  After  the  evolution  of  gas  has  ceased,  re- 
move the  outlet  tube  from  the  flask  and  provide  a 


IRON  SALTS.  103 

tightly  fitting  cork.  Boil  down  the  residual  solu- 
tion to  about  one  half,  cork  the  flask  while  still 
full  of  steam,  set  aside  and  allow  to  cool.  Exam- 
ine the  crystals  which  form.  They  are  ferrous 
sulphate.  Interpret  the  phenomena  observed. 
What  is  the  source  of  the  hydrogen?  Compare 
Exs.  47  and  48.  As  evidence  bearing  on  this  point, 
it  should  here  be  stated  that  iron  forms  with  oxy- 
gen an  oxide  containing  less  oxygen  than  that 
obtained  in  Ex.  56  (£),  but  so  difficult  to  prepare 
and  keep  as  to  be  unsuited  for  class  experiments. 
This  oxide  dissolves  in  dilute  sulphuric  acid  with- 
out evolution  of  hydrogen,  yielding  the  same 
ferrous  sulphate.  In  this  connection,  and  as  a 
step  towards  the  answer  to  the  above  question,  let 
the  student  review  the  relations  of  magnesium, 
magnesium  oxide,  and .  magnesium  hydrate  to- 
wards sulphuric  acid. 

Ex.  58.  Two  Classes  of  Iron  Salts.—  It  is  by 
no  means  true  of  all  metallic  oxides  that  they  will 
dissolve  directly  in  acids  to  form  salts  without 
the  intervention  of  any  other  agent  or  the  forma- 
tion of  any  other  product.  The  oxides  that  do 
sustain  this  relation  to  acids,  like  most  of  those 
we  have  studied,  have  been  called  for  distinction 
sake  salifiable  oxides.  Most  of  the  metals,  like 
sodium,  magnesium,  and  zinc,  form  but  one  sali- 
fiable oxide ;  but  there  are  a  number  of  metals 
which  yield  two  such,  and  it  is  a  remarkable  fact 


104:  LABORATORY  PRACTICE. 

that  the  two  classes  of  salts  thus  formed  differ 
widely  from  each  other  in  their  properties  and  re 
lations.  Iron  is  an  example  in  point,  and  the  two 
classes  of  salts  thus  formed  are  distinguished  as 
ferrous  and  ferric  salts.  The  green  transparent 
crystals  (green  vitriol  or  ferrous  sulphate)  formed 
by  dissolving  iron  in  dilute  sulphuric  acid  is  a 
ferrous  salt,  and  the  same  product,  as  has  been 
said,  also  results  when  the  above-mentioned  fer- 
rous oxide  is  dissolved  in  the  same  acid.  The 
oxide  formed  in  Ex.  56  (&),  called  ferric  oxide,  is 
also  salifiable,  but  when  once  ignited,  as  in  that 
experiment,  dissolves  in  acids  with  difficulty.  The 
hydrate  formed  by  slow  oxidation  in  the  air  in 
contact  with  water  dissolves  readily.  To  substan- 
tiate that  the  ferric  salts  thus  formed  are  essen- 
tially different  from  the  ferrous  salts,  dissolve  a 
portion  of  the  residue  from  Ex.  56  (c)  in  dilute 
sulphuric  acid,  evaporate  nearly  to  dryness  on  a 
watch  glass,  and  try  to  crystallize  the  residue. 
Compare  this  ferric  sulphate  with  the  ferrous  sul- 
phate before  made. 

Ex.  59.  (a)  Iron  and  Sulphur.—  Thoroughly 
mix  5  grammes  of  iron  powder  with  2*8  grammes 
of  flowers  of  sulphur.  Save  a  small  portion  of  the 
mixture  for  comparison ;  heat  the  rest  in  a  small 
flask  (150  cubic  centimetres)  until  the  mass  glows. 
After  cooling  remove  with  a  rod  a  few  grains  of 
the  product  and  compare  with  the  mixture.  (Use 


HYDROGEN  SULPHIDE.  105 

microscope  and  magnet.)     The  black  product  is 
called  iron  sulphide. 

(b)  Hydrogen  Sulphide.—  Leaving  residue  in 
the  flask,  to  which  has  been  fitted  cork  and  out- 
let tube,  pour  in  10  grammes  of  strong  sulphuric 
acid  mixed  with  50  grammes  of  water.  After  air 
has  been  driven  from  the  flask,  collect  the  first 
portions  of  the  escaping  gas  in  large  test  tubes 
over  hot  water.  Ignite  the  gas  at  the  open  mouth 
of  these  test  tubes,  and  observe  the  colour  and 
odour  produced  by  the  flame.  Let  the  rest  of  the 
escaping  gas  bubble  up  through  cold  water  in  a 
glass  -  stoppered  bottle,  and  when  action  is  ex- 
hausted withdraw  outlet  tube,  stopper  the  bottle, 
and  set  aside  for  future  use.  Boil  now  contents  of 
flask  as  in  Ex.  57  (&),  and  set  aside  to  crystallize. 
These  crystals  will  at  once  be  recognized  as  fer- 
rous sulphate,  and  the  nauseous  -  smelling  gas, 
which  is,  to  a  limited  extent,  soluble  in  water,  is 
called  hydrogen  sulphide.  It  is  one  of  the  most 
important  of  chemical  reagents.  As  this  gas  is 
not  only  nauseous,  but  also  to  some  extent  poi- 
sonous, this  experiment  must  be  made  under  a 
hood  or  in  the  open  air.  Interpret  all  the  phases. 

(1)  Iron  sulphide  obviously  consists  of  iron  and  sulphur. 
Iron  dissolved  in  dilute  sulphuric  acid  yields  ferrous  sulphate 
and  hydrogen  gas.  Ferrous  oxide  dissolved  in  the  same  acid 
yields  also  ferrous  sulphate,  but  no  free  gas,  because  the  hy- 
drogen, otherwise  formed,  unites  with  the  oxygen  of  the 


106  LABORATORY  PRACTICE. 

oxide  to  form  water.  Ferrous  sulphide  dissolved  in  the 
same  acid  yields,  again,  ferrous  sulphate  and  a  nauseous- 
smelling  gas.  What  must  be  the  composition  of  this  gas  ? 
Do  the  phenomena  observed  when  the  gas  burns  confirm 
your  inference  ? 


CHAPTER  II. 

GENEEAL  PRINCIPLES. 

13.  Province  "of  Chemistry. 

AT  this  stage  of  his  study  the  student,  having 
become  familiar  with  the  distinctions  implied  by 
the  word  "  substance,"  and  having  acquired  some 
knowledge  of  chemical  phenomena,  is  prepared 
to  understand  what  is  the  province  of  the  science 
of  chemistry.  Chemistry  comprises  and  classi- 
fies our  knowledge  of  those  phenomena  which 
imply  a  change  of  substance.  The  science  of 
physics,  on  the  other  hand,  deals  with  phenomena 
which  do  not  necessarily  imply  a  change  of  sub- 
stance ;  and  hence  the  distinction  between  chemi- 
cal and  physical  changes.  This  distinction  should 
be  illustrated  by  the  teacher  from  the  experiments 
already  made  by  the  student. 

In  every  chemical  change  one  or  more  sub- 
stances, called  the  factors,  change  into  one  or 
more  other  substances,  called  the  products;  and 
it  is  a  primary  object  in  the  study  of  chemistry  to 
learn  what  are  the  factors  and  what  are  the  prod- 
ucts of  every  process  that  comes  under  notice. 


108  LABORATORY   PRACTICE. 

Substances  may  be  mixed  with  one  another,  like 
the  ingredients  of  gunpowder,  or  one  substance 
may  be  dissolved  in  another,  like  salt  in  water, 
without  undergoing  chemical  change ;  but  in  all 
such  cases  the  qualities  of  the  original  substances 
may  be  recognized  in  the  mixture  or  solution. 
Hence  the  very  broad  distinction  between  a  mixt- 
ure, or  a  solution,  and  a  chemical  combination, 
which  the  following  experiments  will  illustrate. 

Ex.  60.  Mixture  and  Chemical  Compound.— 
Mix  together  in  a  mortar  as  intimately  as  possible 
3 '26  grammes  of  zinc  dust  with  1*60  gramme  of 
flowers  of  sulphur.  First  examine  a  small  amount 
of  this  powder  under  a  microscope  of  sufficient 
power  to  show  the  yellow  grains  of  sulphur  lying 
side  by  side  with  the  metallic  grains  of  zinc. 
Make  with  the  rest  of  the  mixture  a  conical  pile 
on  a  square  of  asbestos  paper,  and  apply  the  flame 
of  a  match.  A  chemical  change  ensues,  marked  by 
a  brilliant  deflagration  ;  and  as  the  result  there  is 
left  on  the  paper  a  white  powder,  which  is  a  com- 
pound of  zinc  with  sulphur,  and  is  called  sulphide 
of  zinc.  Here  there  were  two  factors  of  the  chem- 
ical change,  zinc  and  sulphur,  and  one  product,  sul- 
phide of  zinc.  Examine  with  the  microscope  this 
product,  and  no  traces  can  be  seen  either  of  zinc  or 
of  sulphur.  The  deflagration  was  a  manifestation 
of  the  heat  evolved  by  the  chemical  change ;  and 
in  every  chemical  change  there  is  either  a  setting 


PHYSICAL  AND  CHEMICAL  SOLUTION.          109 

free  or  an  absorption  of  energy,  usually  as  heat. 
But  this  last  feature  of  a  chemical  process  will  be 
considered  later  by  itself.  Compare  Ex.  59  (a). 

Ex.  61.  Physical  and  Chemical  Solution.  — 
Take  two  portions  of  one  gramme  each  of  sodic 
carbonate.  Dissolve  one  portion  in  three  cubic 
centimetres  of  water,  evaporate  to  dryness  slowly, 
and  compare  the  residue  with  the  original  salt  in 
appearance,  crystalline  form,  and  taste.  Dissolve 
the  second  portion  in  dilute  hydrochloric  acid, 
evaporate,  and  compare.  What  are  the  factors 
and  what  are  the  products  of  the  chemical  change 
in  the  second  case?  Notice  that  water  is  the 
medium  of  the  chemical  process,  and  dissolves 
one  of  the  products ;  so  that  we  have  here  both  a 
chemical  change  and  a  simple  solution.  The  same 
is  true  when  zinc  dissolves  in  dilute  sulphuric 
acid  (E.v.  47),  or  when  copper  dissolves  in  dilute 
nitric  acid — Ex.  41  (a) — and  the  double  use  of  the 
term  "  solution  "  must  be  made  clear.  What  are 
the  factors  and  what  are  the  products  in  the  two 
cases  of  chemical  solution  last  cited  ?  In  the  same 
way  the  teacher  should  review  the  experiments 
which  thp  strident  has  made,  and  point  out  what 
are  the  factors  and  what  are  the  products  in  each 
case. 


110  LABORATORY  PRACTICE. 

14.   Fundamental  Laws. 

In  every  well-marked  chemical  change  three 
fundamental  laws  are  observed,  and  these  are 
called  the  law  of  conservation  of  mass,  the  law  of 
definite  proportions  by  weight,  and  the  law  of 
definite  proportions  by  volume. 

Ex.  62.  Law  of  Conservation  of  Mass. — Re- 
peat Ex.  24,  but  after  adjusting  the  apparatus 
with  the  bit  of  phosphorus  in  the  spoon  and  fast- 
ening the  cover,  balance  the  jar  on  the  pan  of  the 
balance  with  a  second  jar  of  the  same  volume  and 
such  additional  tare  as  may  be  needed.  Ignite 
now  the  phosphorus  with  a  burning-glass,  and 
after  the  chemical  action,  when  the  jar  is  cold, 
replace  it  on  the  balance.  If  the  jar^-was  tight 
there  will  have  been  no  change  of  weight.  ^IJejpLce 
it  must  be  that— 

The  sum  of  the  weigTits  of  the:  products  of  a 
chemical  change  is  exactly  equal  to  the  sum  of 
the  weights  of  the  factors.  "*•*«•*«--  - 

We  may  conceive  of  any  chemical  process  as 
taking  place  in  an  hermetically  sealed  space — in- 
deed, the  earth  is  essentially  such  a  space — and 
hence  this  law  must  be  universally  true.  The  re- 
sult of  this,  experiment  might  be  anticipated,  and 
it  may  therefore  be  thought  unnecessary ; .  but  its 
very  form  will  make  evident  to  the  student  that 
the  law  of  conservation  of  mass  is  in  .harmony 


FUNDAMENTAL  LAWS.  HI 

with  general  principles  which  he  already  recog- 
nizes. 

Ex.  63.  Law  of  Definite  Proportions  by 
Weight. — Take  five  grammes  of  sal-soda  (crystal- 
lized sodic  carbonate),  selecting  material  that  has 
not  effloresced ;  dissolve  in  dilute  hydrochloric 
acid,  as  directed  in  Ex.  61,  taking  care  to  avoid 
loss  during  the  effervescence  ;  evaporate  to  dry- 
ness,  and  weigh  the  residual  salt ;  calculate  the 
ratio  of  the  sal-soda  used  to  the  salt  produced; 
repeat  the  same  determination  with  ten  grammes 
of  sal-soda,  and  within  the  limits  of  experimental 
error  the  ratio  will  be  the  same  as  obtained  be- 
fore, and  so  would  it  be  whatever  the  amount  of 
sal- soda  employed.  In  this  chemical  change  the 
factors  are  sal-soda  and  hydrochloric  acid,  while 
the  products  are  common  salt,  carbonic  dioxide, 
and  water.  The  last  two,  being  volatile,  escape 
during  the  effervescence  and  subsequent  evapora- 
tion. By  this  experiment  we  have  proved  that 
the  proportion  between  the  weight  of  the  sal-soda 
and  the  weight  of  the  common  salt  is  definite,  and 
it  could  readily  be  shown  experimentally  that  the 
proportion  between  any  two  of  the  five  sub- 
stances involved  in  this  chemical  change  was 
equally  definite.  So  of  any  other  well-marked 
chemical  change,  and  hence  the  general  law 
that — 

In  any  well-marked  chemical  change  the  rela- 


112  LABORATORY  PRACTICE. 

tive  weights  of  the  several  factors  and  products 
are  definite  and  invariable. 

Here,  again,  the  result  might  have  been  antici- 
pated, for  it  only  amounts  to  finding  that  if  we 
use  twice  as  much  sal-soda  we  shall  obtain  twice 
as  much  common  salt,  which  might  seem  self- 
evident;  and  this  consideration  will  show  that 
the  law  of  definite  proportions  by  weight  is  in 
entire  harmony  with  principles  universally  rec- 
ognized. 

Ex.  64.  Law  of  Definite  Proportions  by  Vol- 
ume.— This  law,  sometimes  called  the  law  of  Gay- 
Lussac,  may  be  thus  stated : 

In  any  well-marked  chemical  change  the  rela- 
tive volumes  of  the  aeriform  factors  or  products, 
if  measured  under  the  same  conditions,  bear  to 
each  other  a  simple  numerical  ratio. 

It  has  already  been  illustrated  by  several  ex- 
periments, which  it  is  unnecessary  to  repeat. 
Thus  it  was  shown  by  Ex.  32  that  when  oxygen 
combines  with  sulphur  to  form  sulphurous  oxide 
the  volume  of  this  sole  product  is  the  same  as  the 
volume  of  the  oxygen  gas  used.  A  similar  rela- 
tion appeared  when  oxygen  united  with  carbon 
to  form  carbonic  dioxide  in  Ex.  24  (c).  Again, 
when,  in  Ex.  40  (a),  carbonic  dioxide  united  with 
more  carbon  to  form  carbonic  oxide  the  volume 
of  the  gas  was  doubled.  A  still  more  striking  il- 
lustration of  the  law  is  to  be  found  in  the  fact 


ELEMENTARY  SUBSTANCES.  H3 

that  two  volumes  of  hydrogen  gas  combine  with 
one  volume  of  oxygen  gas  to  form  two  volumes  of 
vapour  of  water,  all  measured,  of  course,  under 
the  same  pressure  and  at  a  temperature  above  the 
boiling  point  of  water.  The  experiment  is  easily 
made  with  a  form  of  eudiometer  invented  for  the 
purpose  by  Hofmann  and  sold  by  all  the  dealers 
in  chemical  apparatus,  and  it  should  be  shown  to 
the  class  if  possible.  (_ 


15.   Compounds  and  Elements. 

The  student  can  not  have  performed  the  ex- 
periments heretofore  described  without  himself 
drawing  the  inference  in  certain  cases  that  the 
products  have  been  formed  by  the  union  of  two 
or  more  factors,  and  in  other  cases  that  the  prod- 
ucts have  resulted  from  the  breaking  up  of  a 
factor  into  simpler  parts.  Hence  come  the  fun- 
damental conceptions  of  composition  and  decom- 
position, of  synthesis  and  of  analysis,  as  we  have 
previously  called  them.  Our  judgment  in  any 
case  depends  not  only  on  the  circumstances  of  the 
experiment,  but  also  on  a  comparison  of  the 
weights  of  the  products  with  those  of  the  factors 
from  which  they  were  formed.  Thus,  in  Ex.  24 
(a),  it  is  perfectly  evident  from  the  conditions  of 
the  experiment  that  the  white  product  results 
from  the  union  of  phosphorus  and  oxygen.  If 


LABORATORY  PRACTICE. 

now  in  addition  we  could  weigh  the  white  prod- 
net  and  find  that  its  weight  was  exactly  equal  to 
that  of  the  phosphorus  and  oxygen  used,  the 
proof  of  its  composition  would  be  complete.  So 
also  when,  in  Ex.  16,  we  pass  a  current  of  elec- 
tricity through  water  and  see  oxygen  and  hydro- 
gen gases  escaping  from  the  platinum  poles  of  the 
apparatus,  and  notice  that  everything  else  re- 
mains unchanged,  we  conclude  that  the  two  gases 
must  come  from  the  water  and  are  the  products  of 
its  decomposition  ;  but  we  do  not  have  absolute 
proof  until,  as  in  Ex.  53  (a),  we  pass  hydrogen 
over  oxide  of  copper  and  find  that  the  weight  of 
the  water  formed  is  exactly  equal  to  that  of  the 
hydrogen  and  oxygen  which  have  disappeared. 
In  like  manner,  our  knowledge  of  the  composition 
of  other  substances  is  the  result  of  our  knowledge 
of  chemical  processes,  which  has  been  accumu- 
lated during  long  years  of  study.  As  the  total 
result  of  this  study,  we  may  say  that  while  the 
larger  number  of  substances  which  we  handle  may 
be  decomposed  or  analyzed,  there  are  about  sev- 
enty known  substances  which  can  not  be  broken 
up  into  simpler  parts,  and  these  we  call  ele- 
mentary substances.  An  elementary  substance 
differs  from  other  substances  only  in  this,  that  it 
enters  into  all  chemical  changes  as  a  whole,  and 
we  know  of  no  chemical  process  in  which  it  be- 
comes divided.  It  does,  however,  enter  into 


ELEMENTARY  SUBSTANCES.  H5 

nnion  with  other  substances ;  and,  speaking  in 
general,  we  may  by  the  combination  of  the  ele- 
mentary substances  reproduce  all  the  forms  of 
matter  with  which  we  are  acquainted.  The  sys- 
tematized knowledge  of  the  methods,  whether 
analytical  or  synthetical,  by  which  the  composi- 
tion of  bodies  has  been  determined  is  a  very  im- 
portant branch  of  chemical  science,  known  under 
the  name  of  chemical  analysis  ;  and  the  subject 
is  subdivided  into  qualitative  and  quantitative 
analysis,  according  as  the  object  in  view  is  to  de- 
termine solely  the  nature  or  the  proportion  of  the 
ingredients.  In  either  case  the  analysis  may  be 
either  ultimate  or  approximate.  It  is  ultimate 
when  we  seek  for  the  elementary  substances  of 
which  the  compound  consists.  It  is  approximate 
when  we  look  for  the  simpler  products  (for  the 
most  part  acids  and  metallic  oxides)  into  which 
the  complex  material  may  be  primarily  divided. 
The  following  experiments  will  give  a  general, 
but  necessarily  a  very  imperfect  idea  of  the  man- 
ner in  which  the  results  are  reached  : 

(1)  A  list  of  the  elementary  substances  will  be  found  in 
the  table  at  the  end  of  this  book,  and  this  should  be  care- 
fully examined  by  the  student  in  reference  to  the  substances 
he  has  met  with  in  the  course  of  his  experiments.  Which 
of  these  are  elements  ?  The  student  should  make  a  list  of 
the  elementary  substances  with  which  he  has  become  famil- 
iar. Can  an  elementary  substance  be  told  by  its  external 
characters  ?  Is  there  not  a  class  of  bodies  which  are  uni- 
formly elementary  ? 


116  LABORATORY  PRACTICE. 

16.  Qualitative  Analysis. 

Ex.  65.  Analysis  of  a  Silver  Com.— Dissolve  a 
ten-cent  coin  in  5  cubic  centimetres  of  pure,  strong 
nitric  acid,  diluted  with  its  own  volume  of  water, 
Dilute  to  50  cubic  centimetres.  Add  hydrochloric 
acid  to  hot  solution  so  long  as  a  precipitate  is  pro- 
duced. Filter,  wash  thoroughly  (three  times) 
with  water,  dry  precipitate.  Transfer  to  glass 
combustion  tube  connected  with  hydrogen  genera- 
tor as  in  Ex.  53  (a).  Interpose  chloride-of -calci- 
um tube  between  generator  and  combustion  tube. 
Allow  hydrogen  to  flow  through  the  apparatus 
long  enough  to  expel  the  air ;  then  heat  the  com- 
bustion tube,  and  continue  until  reduction  is  com- 
plete. Test  gas  evolved  from  outlet,  which  should 
be  bent  downwards.  Preserve  silver.  Add  a  few 
drops  of  sulphuric  acid  to  blue  filtrate  and  evapo- 
rate (under  hood)  to  get  rid  of  the  volatile  acids ; 
dilute  to  10  cubic  centimetres  and  insert  a  strip  of 

zinc. 

(1)  What  is  the  gas  evolved  during  the  reduction? 
What  is  its  composition  ?  What  must  be  the  composition  of 
the  precipitate  ?  Does  the  formation  of  this  precipitate  con- 
form to  any  general  principle — Ex.  14  (3).  On  what  circum- 
stances does  the  separation  of  copper  form  silver  in  this  ex- 
periment ?  How  can  you  be  sure  that  the  copper  obtained 
came  from  the  coin  and  not  from  any  of  the  accessory  ma- 
terials employed  ? 

Ex.  66.  (a)  Analysis  of  Marble.—  Heat  two 
grammes  of  marble  dust  in  a  small  iron  crucible 


ANALYSIS  OP  MARBLE.  H7 

over  a  blast  lamp,  so  long  as  the  material  contin- 
ues to  lose  in  weight.  The  residue  easily  recog 
nized  as  lime  must  be  one  of  the  constituents 
Add  water,  test  with  litmus  paper,  and  compare 
with  a  mixture  of  marble  dust  and  water.  What 
inference  would  you  draw  from  Ex.  48  in  regard 
to  the  probable  constitution  of  such  white  pow- 
ders? We  know  that  magnesium  has  a  very 
strong  attraction  for  oxygen  (Ex.  48),  and  there- 
fore, to  test  this  inference  heat  over  a  lamp  in  a 
small  ignition  tube  two  decigrammes  of  lime 
in  powder  mixed  with  one  decigramme  of  mag- 
nesium powder.  When  cold  add  water,  and. 
shake  up  with  residue.  Note  the  evolution  of 
hydrogen  gas  and  the  production  of  lime  water. 
The  metal  liberated,  which  decomposes  water,  is 
calcium.  Hence  lime  consists  of  calcium  and 
oxygen. 

What  was  driven  off  from  the  marble  dust  ~by 
heat  f  To  show  this,  let  the  teacher  procure  half 
a  metre  of  small  iron  gas  pipe.  Close  one  end  by 
welding.  Drop  in  ten  grammes  of  marble  dust 
and  shake  down  to  the  closed  end.  Mount  in  a 
Fletcher  gas  furnace  and  connect  the  open  end 
with  a  pneumatic  trough.  Heat  to  a  full  white 
heat,  avoid  excess  of  illuminating  gas  at  the 
burner  lest  it  diffuse  through  the  tube.  Col- 
lect and  examine  gas  evolved.  It  will  not  sup- 
port combustion,  it  dissolves  in  water,  and  feebly 


118  LABORATORY  PRACTICE. 

reddens  litmus  paper.  It  is  obviously  carbonic 
dioxide.  Of  what  does  carbonic  dioxide  consist  ? 
Take  a  length  of  magnesium  ribbon.  Ignite  and 
plunge  the  burning  end  in  a  jar  of  carbonic  diox- 
ide. Observe  the  separation  of  carbon.  What  is 
the  necessary  inference  in  regard  to  the  composi- 
tion of  carbonic  dioxide  ?  Discuss  all  points  of 
this  evidence. 

(b)  Synthetical  Confirmation. — Review  in  this 
connection  Ex.  39  (&),  and  repeat  on  a  small  scale 
with  the  lime  water  obtained  by  the  action  of  cal- 
cium on  water.  Discuss  the  phenomena  as  syn- 
thetical evidence  of  the  composition  of  marble. 
Explain  the  action  of  hydrochloric  acid  on  marble, 
bringing  in  contrast  the  two  facts — 

Marble  and  hydrochloric  acid  yields  calcic 
chloride  and  carbonic  dioxide. 

Lime  and  hydrochloric  acid  yield  calcic  chlo- 
ride. 

Confirm  these  facts  experimentally  and  draw 
your  own  inferences.  As  will  hereafter  appear, 
the  facts  as  above  stated  are  not  complete  state- 
ments, since  in  both  processes  water  is  also  formed, 
which  in  qualitative  experiments  escapes  notice 
by  mixing  with  the  mass  of  liquid,  acting  as  the 
medium  of  the  chemical  change.  Still  in  the  pres- 
ent case  the  fact  overlooked  was  not  material, 
and,  since  in  experimental  science  we  can  fre- 
quently draw  correct  conclusions  from  similar  in 


CHEMICAL  TESTS.  119 

complete  evidence,  this  experience  may  teach  a 
valuable  lesson.  It  was  from  exactly  this  evi- 
dence that  the  proximate  composition  of  marble 
was  first  inferred  by  Dr.  Black,  of  Edinburgh,  a 
century  ago. 

Synthetical  processes  are  often  of  great  value 
in  confirming  analytical  results,  and  give  fre- 
quently the  most  direct  and  efficient  means  of 
finding  out  the  composition  of  a  material.  As 
commonly  used  the  term  chemical  analysis  in- 
cludes all  methods  of  establishing  the  chemical 
constitution  of  substances,  whether  synthetical  or 
strictly  analytical. 

Ex.  67.  Chemical  Tests. — In  the  examples  of 
analysis  given  above,  we  have  actually  separated 
the  elementary  substances  of  which  two  familiar 
bodies  consist.  But  such  a  separation  is  not  al- 
ways practicable  or  necessary,  and  we  can  gen- 
erally discover  the  constituents  of  a  substance  by 
applying  certain  characteristic  tests. 

Take  four  short  lengths,  not  over  50  millimetres 
long,  of  platinum  wire  and  make  a  loop  at  the  end 
of  each  ;  melt  into  the  first  of  these  loops  chloride 
of  sodium  ;  into  the  second,  chloride  of  potassium  ; 
into  the  third,  chloride  of  strontium  ;  and  into  the 
fourth,  chloride  of  barium.  Hold  the  loops  suc- 
cessively in  the  flame  of  a  Bunsen  lamp,  and  no- 
tice the  colours  which  they  impart  to  it ;  and  if  a 
spectroscope  is  accessible,  examine  the  coloured 


120  LABORATORY  PRACTICE. 

flames  with  this  instrument.  Almost  any  prepa- 
ration of  sodium,  potassium,  strontium,  or  bari- 
um would  produce  the  same  effect ;  and  these 
characteristic  colors,  or  still  better  the  correspond- 
ing bands  seen  in  the  spectroscope,  are  indica- 
tions, or,  as  we  usually  say,  tests  of  these  metals. 
For  another  illustration,  take  five  test  tubes ;  in 
the  first  dissolve  a  small  amount  of  zinc  dust  in 
acetic  acid ;  in  the  second,  a  bit  of  white  arsenic 
in  hydrochloric  acid  ;  in  the  third,  a  bit  of  anti- 
mony in  hydrochloric  acid  to  which  has  been 
added  a  drop  of  nitric  acid  ;  in  the  fourth,  a  bit  of 
iron  in  hydrochloric  acid  ;  in  the  last,  a  bit  of  lead 
in  weak  nitric  acid.  In  each  case  use  a  bit  of 
metal  no  larger  than  a  pin's  head,  and  dissolve  in 
the  least  amount  of  acid  possible  ;  half  fill  the 
test  tubes  with  water ;  add  to  the  first  three  the 
solution  of  sulphide  of  hydrogen  obtained  in  Ex. 
59  (b) ;  to  the  fourth,  a  few  drops  of  a  solution  of 
ferrocyanide  of  potassium  ;  and  to  the  last,  a  few 
drops  of  a  solution  of  potassic  chromate.  The 
characteristically  coloured  precipitates  obtained 
under  the  conditions  present  are  in  each  case 
tests  of  the  several  metals.  So,  in  general,  it  is 
not  necessary  to  isolate  an  acid,  a  metallic  oxide, 
or  an  elementary  substance,  in  order  to  prove 
that  it  is  present  or  absent  in  a  given  case,  but 
only  to  use  the  proper  tests  in  the  right  way ; 
and  the  works  on  qualitative  analysis  teach  us 


QUANTITATIVE  ANALYSIS.  121 

in  what  order  and  under  what  conditions  the 
proper  tests  should  be  applied.  The  practice  of 
qualitative  analysis  affords  a  most  admirable  train- 
ing in  the  methods  of  inductive  reasoning. 


17.   Quantitative  Analysis. 

Ex.  68.  Analysis  of  Potassium  Bromide. — 
The  analysis  made  in  Ex.  65  may  be  made  quanti- 
tative by  first  weighing  the  coin  and  afterwards 
weighing  the  silver  and  copper  obtained.  Of 
course  if  the  coin  consists  of  nothing  else,  the  sum 
of  the  weights  of  the  two  metals  ought  to  exactly 
equal  the  weight  of  the  coin,  and  such  a  coinci- 
dence would  go  far  to  establish  the  accuracy  of 
our  work.  In  the  analysis  as  above  made  only  a 
very  rough  approximation  to  equality  could  be 
expected,  but  by  more  accurate  methods  such  a 
confirmation  of  the  work  could  be  almost  abso- 
lutely secured.  In  general,  however,  in  order  to 
determine  the  relative  proportions  of  the  different 
constituents  in  a  compound,  it  is  rarely  practica- 
ble to  separate  the  ingredients  and  weigh  the  sev- 
eral amounts.  The  method  is  to  transfer  each 
ingredient  to  some  new  combination  which  can  be 
formed  without  loss,  weighed  with  accuracy,  and 
the  composition  of  which  through  previous  an- 
alyses is  absolutely  known.  Take  the  simplest 
case.  We  wish  to  analyze  common  salt,  which  is 


122  LABORATORY   PRACTICE. 

known  to  consist  wholly  of  chlorine  and  sodinm. 
Neither  of  these  are  ingredients  which  can  be  ac- 
curately separated  and  weighed.  So,  with  a  care- 
fully weighed  quantity  of  salt,  we  prepare  chlo- 
ride of  silver  by  a  process  in  which  we  are  sure 
that  every  particle  of  chlorine  has  been  trans- 
ferred from  its  previous  combination  with  sodium 
to  a  new  combination  with  silver.  Chloride  of 
silver  is  a  substance  which  can  be  collected  and 
weighed  with  perfect  precision.  It  has  previously 
been  accurately  analyzed  over  and  over  again,  and 
by  referring  to  tables  we  find  what  fraction  of  the 
weight  of  chloride  of  silver  thus  found  consists  of 
chlorine,  and  then  a  very  simple  calculation  gives 
the  weight  of  chlorine  sought.  If  the  salt  is  abso- 
lutely pure  the  rest  of  the  weight  taken  consists 
of  sodium,  and  the  amount  of  sodium  could  not 
be  determined  so  accurately  in  any  other  way. 
Obviously,  the  analysis  of  common  salt  rests  back 
on  the  analyses  of  chloride  of  silver  previously 
made  and  recorded  ;  and  so  in  most  cases  our  an- 
alyses of  to-day  rests  back  on  the  work  of  those 
who  have  gone  before  us.  After  these  relations 
have  been  explained,  let  the  student  weigh  in  a 
small  beaker  glass  exactly  one  gramme  of  pure 
potassic  bromide,*  and  dissolve  the  salt  in  about 

*  With  the  rude  manipulation  here  expected,  potassium  bromide 
will  give  more  precise  results  than  common  salt  and  illustrates 
equally  well  the  general  methods  of  quantitative  analysis.  Potas- 


QUANTITATIVE  ANALYSIS.  123 

twenty-five  cubic  centimetres  of  water.  Weigh  in 
a  similar  beaker  one  and  a  half  gramme  of  silver 
nitrate  and  dissolve  in  an  equal  amount  of  water. 
Pour  now  with  constant  stirring  the  first  solution 
into  the  second,  rinse  the  beaker,  wash  in  the  last 
drops,  and  allow  to  stand  until  the  precipitate 
fully  settles.  Collect  on  a  tared  filter,  wash  dry, 
and  weigh.  It  is  known  from  previous  work,  and 
can  be  found  by  reference  to  any  work  on  quanti- 
tative analysis,  that  every  gramme  of  silver  bro- 
mide contains  0*4255  gramme  of  bromine ;  and 
practically  an  analyst  would  always  assume  that 
this  value  was  given  and  at  once  calculate  the 
amount  of  bromine  in  the  weight  of  the  silver  bro- 
mide he  had  obtained,  and  this  would  be  the 
amount  in  the  weight  of  the  bromide  of  potassium 
he  had  taken  for  analysis.  To  show  the  student, 
however,  that  results  thus  obtained  rest  back  on 
previous  analyses,  let  him  make  this  additional 
determination,  not  that  he  can  compete  with  the 
old  work,  which  has  been  often  repeated  with  the 
greatest  care  in  order  to  establish  fundamental 
data  for  just  such  uses  as  are  here  indicated,  but 
to  the  end  that  he  may  realize  the  actual  relations 
in  most  problems  of  quanitative  analysis.  Weigh 
out  exactly  one  gramme  of  pure  metallic  silver, 

slum  bromide  is  a  familiar  medicinal  salt,  consisting  of  potassium 
and  bromine,  two  elementary  substances  closely  allied  respectively 
to  sodium  and  chlorine. 


124  LABORATORY  PRACTICE. 

place  in  a  small  beaker,  .and  dissolve  in  about  two 
cubic  centimetres  of  strong  nitric  acid  diluted 
with  four  cubic  centimetres  of  water,  and  add 
twenty-five  cubic  centimetres  of  water.  In  a  sec- 
ond beaker  dissolve  in  twenty-five  cubic  centime- 
tres of  water  1-2  gramme  of  potassic  bromide, 
and  then  proceed  as  before.  From  the  result  cal- 
culate the  weight  of  bromine  in  one  gramme  of 
silver  bromide,  which  should  be,  within  the  limits 
of  error,  the  same  as  the  value  given  above.  The 
student  should  now  calculate  from  his  own  results 
the  per  cent  of  bromine  and  of  potassium  in  po- 
tassic bromide,  and  should  have  practice  in  simi- 
lar calculations  until  he  is  familiar  with  the  usual 
manner  of  stating  the  results  of  analysis  in  per 
cent. 

(1)  It  is  known  that  pure  common  salt  consists  wholly 
of  sodium  and  chlorine,  and  also  that  one  gramme  of  silver 
chloride  contains  0'2474  gramme  of  chlorine.    In  one  deter- 
mination 0*5723  gramme  of  salt  gave  1"4038  gramme  of  sil-  . 
ver  chloride.     Calculate  the  percentage  composition  of  com- 
mon salt. 

Ans.  Chlorine,    60 '69 
Sodium,     39-31 

100-00 

(2)  It  is  known  that  pure  crystallized  cane  sugar  con- 
sists wholly  of  carbon  hydrogen  and  oxygen.     By  the  usual 
process  of  organic  analysis,  0'2569  gramme  of  sugar  gave 
0'3966  gramme  of  carbonic  dioxide  and  0"1487  gramme  of 
water.     It  is  known  that  one  gramme  of  carbonic  dioxide 
contains  0*2727  gramme  of  carbon,  and  one  gramme  of  water 
0*1111  gramme  of  hydrogen.     What  is  the  percentage  com- 
position of  cane  sugar  ? 


QUANTITATIVE  ANALYSIS.  125 

Ans.  Carbon,  42 '11 
Hydrogen,  6 '43 
Oxygen,  51 '46 

100-00 

(3)  Obviously  the  above  processes  assume  a  qualitative 
knowledge  of  the  composition  of  the  substance  analyzed, 
and  so,  in  general,  quantitative  analysis  implies  a  previous 
qualitative  analysis.  Indeed,  the  process  of  determining  the 
amount  of  an  ingredient  present  must  constantly  be  varied 
according  as  it  is  associated  with  different  substances,  and  a 
large  knowledge  is  required  in  order  to  meet  the  conditions 
in  any  case  and  secure  accurate  results.  Thus  quantitative 
analysis  becomes  a  distinct  and  widely  extended  branch  of 
chemical  study,  and  it  is  the  chief  work  of  the  practical 
chemist. 


CHAPTER  III. 

MOLECULES  AND   ATOMS. 

18.  Molecular  Theory. 

THE  theory  of  the  new  science  of  thermo-dy- 
namics  assumes  that  the  material  of  aeriform  bod- 
ies is  not  continuously  distributed  through  the 
spaces  they  seem  to  fill,  but  consists  of  a  vast 
number  of  exceedingly  minute  particles  in  rapid 
motion  to  and  fro,  constantly  rebounding  from 
one  another,  or  from  the  walls  of  the  containing 
vessel.  These  minute  particles  are  called  mole- 
cules, and  the  phenomena  of  heat  are  supposed 
to  be  manifestations  of  their  moving  power.  The 
molecules  of  the  same  substance,  of  water,  for  ex- 
ample, are  supposed  to  have  the  same  weight — in 
fact,  to  be  alike  in  every  respect ;  while  the  mole- 
cules of  different  substances  are  as  unlike  as  the 
substances  themselves.  This  theory  has  been 
worked  out  mathematically  with  great  ability, 
and  the  phenomena  of  nature  have  been  found,  in 
a  most  remarkable  manner,  to  conform  to  the  de- 
ductions of  mathematical  analysis.  Of  these  de- 
ductions one  of  the  most  remarkable  is,  that — 


MOLECULAR  WEIGHT.  127 

Equal  volumes  of  all  gases  or  vapours,  meas- 
ured under  the  same  conditions,  contain  the  same 
number  of  molecules. 

This  deduction  is  usually  called  the  law  of 
Avogadro  ;  and  if  we  accept  the  fundamental  con- 
ception of  molecular  structure  we  must  also  ac- 
cept this  inference  which  it  involves.  The  mod- 
ern theory  of  chemistry  accepts  the  law  of  Avo- 
gadro as  a  fundamental  principle,  and  builds 
upon  it  a  large  superstructure. 

The  law  of  Avogadro  does  not  absolutely  hold 
except  when  the  material  is  in  a  perfectly  aeri- 
form condition.  It  is  only  approximately  true  in 
the  case  of  dense  gases  or  vapours  under  the 
pressure  of  the  air,  and  near  the  point  of  con- 
densation the  deviations  are  sometimes  very 
marked.  It  has  no  reference  whatever  to  liquids 
or  solids.  These  forms  of  matter  are  supposed 
also  to  consist  of  moving  particles  ;  but,  if  so,  the 
molecules  must  be  variously  compacted,  and  their 
motions  otherwise  circumscribed  than  in  the  aeri- 
form state. 

19.  Physical  Method  of  Determining  Molec- 
ular Weights. 

If  the  law  of  Avogadro  is  true,  the  molecular 
weight  of  a  substance  must  be  proportioned  to  its 
specific  gravity  in  the  state  of  gas  or  vapour.  If 


128  LABORATORY  PRACTICE. 

we  take  hydrogen  gas  as  the  unit  of  reference  for 
the  specific  gravity,  and  the  molecule  of  hydrogen 
gas  as  the  unit  of  reference  for  molecular  weights, 
then  the  number  which  expresses  the  specific 
gravity  of  a  substance  in  the  state  of  gas  or  vapour 
would  also  express  the  molecular  weight  of  that 
substance.  For  considerations  which  will  shortly 
appear,  half  the  weight  of  the  molecule  of  hydro- 
gen has  been  taken  as  the  unit  of  molecular 
weight,  so  that  the  molecule  of  hydrogen  gas 
weighs  two  of  the  assumed  units ;  and  hence,  on 
this  system,  the  molecular  weight  of  any  substance 
is  found  by  doubling  its  specific  gravity  taken  in 
the  aeriform  state,  and  referred  to  hydrogen  gas 
as  the  standard. 

The  physical  method  of  determining  molecular 
weights,  therefore,  reduces  itself  to  finding  the 
specific  gravity  of  a  substance  when  in  the  condi- 
tion, of  gas  or  vapour  with  reference  to  hydrogen. 
The  substance  must  be  in  the  condition  of  gas  or 
vapour,  and  the  method  is  only  applicable  to  those 
bodies  which  are  naturally  aeriform,  or  which  can 
be  volatilized  at  temperatures  within  control 
without  undergoing  decomposition. 

Ex.  69.  Density  of  Hydrogen. — Use  a  flask 
not  exceeding  100  cubic  centimetres  capacity. 
Cork  tightly,  and  connect  through  cork  a  small 
chloride  of  calcium  tube,  so  proportioning  the 
parts  that  the  flask  will  stand  on  the  pan  of  the 


MOLECULAR  WEIGHT.  129 

balance.  Add  to  the  flask  20  cubic  centimetres  of 
strong  hydrochloric  acid  and  an  equal  volume  of 
water.  Weigh  out  closely  5  grammes  of  sheet 
zinc  that  has  been  carefully  cleaned.  Place  this 
at  the  side  of  the  flask  on  the  scale  pan  and  take 
the  tare  as  closely  as  possible.  Kemoving  now 
the  flask  from  the  balance,  connect,  by  means  of  a 
flexible-rubber  connector,  the  chloride- of -calcium 
tube  with  the  inlet  tube  of  the  gasometer  before 
described  (note  to  Ex.  45),  which  should  be  large 
enough  to  hold  two  full  litres  of  gas.  When  all 
is  ready  withdraw  the  cork,  drop  in  the  zinc,  and 
quickly  recork  the  flask.  Wait  until  the  evolu- 
tion of  hydrogen  has  altogether  ceased,  then  shut 
off  the  gasometer  and  disconnect  the  flask.  Be- 
fore replacing  it  on  the  balance  pan  withdraw 
the  cork  for  a  few  minutes  to  give  the  hydrogen 
gas  in  the  interior  time  to  diffuse  into  the  air. 
The  loss  of  weight  is  obviously  the  weight  of  the 
hydrogen  set  free.  To  find  the  volume  of  this 
hydrogen  transfer  the  gas  from  the  gasometer  in 
successive  portions  to  a  litre  measure  over  a  pneu- 
matic trough.  In  reading  the  volumes  take  care 
that  the  level  of  the  water  is  the  same  inside  and 
outside  the  glass  and  avoid  warming  with  the 
hands.  This  volume  should  be  corrected  for  the 
tension  of  aqueous  vapour  (Ex.  22)  and  reduced  to 
standard  conditions,  observing,  for  the  purpose, 
the  height  of  the  barometer  and  the  temperature 


130  LABORATORY  PRACTICE. 

of  the  water  in  the  trough.  We  have,  then,  the 
weight  of  a  measured  volume  of  hydrogen,  and 
can  easily  calculate  the  weight  of  one  litre.  With 
such  appliances  as  are  here  assumed  the  process 
is  accurate  within  about  5  per  cent.  On  account 
of  its  great  lightness  the  exact  determination  of 
the  density  of  hydrogen  gas  is  a  difficult  problem, 
and  for  the  purpose  of  calculating  the  specific 
gravity  of  other  aeriform  bodies  referred  to  hy- 
drogen as  unity,  we  will  assume  the  value  gener- 
ally received,  0'0896,  and  for  the  specific  gravity 
referred  to  air,  0'0692.  (Compare  Ex.  17.) 

Ex.  70.  Specific  Gravity  of  Carbonic  Dioxide. 
-Take  a  quart  tin  can,  with  narrow  neck  fitted 
with  selfsealing  stopper,  and,  having  measured  its 
exact  contents,  as  in  Ex.  17,  carefully  clean  and 
dry  it.  Place  it,  open,  on  the  balance-pan,  and 
equipoise  it  with  a  second  can  of  the  same  size  and 
pattern,  tightly  sealed.  Observe  the  thermometer 
and  barometer  at  the  time  the  equilibrium  is  estab- 
lished. It  will  be  obvious  now  that  if  we  calcu- 
late the  weight  of  air  which  the  open  can  holds  at 
this  temperature  and  pressure  (Ex.  17)  and  add 
an  equal  weight  to  the  pan  carrying  the  open  can, 
we  should  have  what  would  be  the  exact  tare  if 
the  can  were  sealed  with  all  the  air  exhausted 
from  the  interior.  Moreover,  since  the  can  and 
its  counterpoise  displace  the  same  volume  of  air, 
it  is  also  obvious  that  this  equilibrium  would  not 


MOLECULAR  WEIGHT.  131 

be  disturbed  by  any  changes  in  the  atmosphere. 
If,  therefore,  we  fill  the  can  with  any  gas — for 
example,  carbonic  dioxide — the  increased  weight 
will  be  simply  the  weight  of  this  gas.  Remove 
then  the  open  can,  fill  it  with  carbonic-dioxide 
gas  by  displacement,  and  seal  it,  observing  the 
thermometer  and  barometer  at  the  moment  the 
can  is  closed.  Determine  the  increased  weight ; 
and  this  is  the  weight  at  the  last  observed  tem- 
perature and  pressure  of  a  volume  of  carbonic  - 
dioxide  gas  equal  to  the  known  capacity  of  the 
can.  From  this  value  calculate  what  would  be 
the  weight  of  the  same  volume  if  the  thermome- 
ter marked  0°  and  the  barometer  stood  at  30 
inches.  The  density  of  carbonic-dioxide  gas  — 
that  is,  the  weight  of  a  litre  under  the  standard 
conditions  of  temperature  and  pressure — is  then 
found  by  dividing  the  weight  of  the  gas  by  the 
capacity  of  the  can.  The  density  of  carbonic 
dioxide  divided  by  the  density  of  air  gives  the 
specific  gravity  of  carbon  dioxide  referred  to  air ; 
or,  if  divided  by  the  density  of  hydrogen,  the  spe- 
cific gravity  referred  to  hydrogen  ;  and  these  val- 
ues are  the  same  for  all  temperatures  and  press- 
ures. Why  ? 

(1)  What  is  the  molecular  weight  of  carbonic  dioxide  ? 

(2)  The  specific  gravity  of  nitrous-oxide  gas  referred  to 
hydrogen  is  22*04.     What  is  its  molecular  weight  ? 

(3)  The  specific  gravity  of  cyanogen  gas  referred  to  hy- 
drogen is  26'06.     What  is  its  molecular  weight  ? 


132  LABORATORY  PRACTICE. 

(4)  The  specific  gravity  of  oxygen  gas  referred  to  hydro- 
gen is  16.  What  is  its  molecular  weight  ? 

Ex.  71.  J^peci/ic  Gravity  of  Vapours. — Find 
the  molecular  weight  of  alcohol,  ether,  chloro- 
form, or  ethylene  bromide,  by  determining  in 
either  case  the  specific  gravity  of  the  vapour  of 
the  substance  by  the  following  method  : 

Determine  the  volume  of  one  of  the  bulbs 
provided  for  the  purpose,*  and,  having  thor- 
oughly dried  both  the  interior  and  the  exteri- 
or surface,  seal  the  shorter  tubulature.  Select  a 
second  bulb  of  approximately  the  same  size,  and, 
having  sealed  both  of  its  tubulatures,  use  it  to 
equipoise  the  first,  completing  the  tare  as  conven- 
ient. Observe  now  the  thermometer  and  barome- 
ter, and  calculate  the  weight  of  the  dry  air  which 
fills  the  open  bulb — Ex.  17  (1).  Add  this  weight  to 
the  pan  holding  the  first  bulb  ;  and,  as  thus  load- 
ed, the  balance  would  be  in  equilibrium  were  the 
glass  vessel  completely  exhausted.  Moreover, 
this  constructive  equilibrium  will  not  be  disturbed 
by  any  atmospheric  changes  (Ex.  70).  Introduce 
now  into  the  first  bulb  about  50  cubic  centimetres 
of  the  volatile  liquid  under  examination.  Hang 
the  bulb  above  the  water  in  an  ordinary  tea-ket- 

*  These  bulbs  hold  about  400  cubic  centimetres,  and  are  blown 
in  a  mould  so  as  to  secure  a  uniform  size.  They  have  a  long,  nar- 
row stem  and  opposite  to  the  stem  a  short  and  still  narrower  tubu- 
lature. 


MOLECULAR  WEIGHT.  133 

tie  sufficiently  capacious  for  the  purpose,  with  the 
longer  tubulature  projecting  through  a  cork  fit- 
ting a  hole  in  the  cover.  Boil  the  water  under 
the  bulb  so  that  the.  steam  surrounding  the  glass 
and  escaping  by  the  nozzle  of  the  kettle  shall 
maintain  a  uniform  temperature  near  100°.  Ob- 
serve this  temperature  with  a  thermometer  pass- 
ing through  a  cork  fitted  to  a  second  hole  in  the 
cover,  and  as  soon  as  the  current  of  vapour  from 
the  bulb  stops  seal  the  tubulature  by  melting 
with  a  blowpipe  the  glass  at  the  tip.  At  the  same 
time  note  the  height  of  the  barometer.  When 
cold,  replace  the  bulb  on  the  balance  and  deter- 
mine the  increased  weight  above  the  equilibrium 
just  described.  This  value  is  the  weight  of  the 
vapour  which  filled  the  bulb  when  in  the  kettle 
at  the  temperature  and  pressure  observed.  Find 
next  what  must  have  been  the  slightly  increased 
volume  of  the  bulb  when  in  the  kettle,  using  the 

formula — 

V  =  V  (1  +  0-000024  x  t°), 

and  calculate  what  would  be  the  weight  of  the 
same  volume  of  dry  air  at  the  temperature  and 
pressure  in  the  kettle.  Then  the  weight  of  the  va- 
pour divided  by  the  weight  of  the  air  gives  the  spe- 
cific gravity  of  the  vapour  referred  to  air,  and  this 
result  multiplied  by  14 -43  gives  the  specific  gravity 
of  the  same  vapour  referred  to  hydrogen  gas.  To 
make  sure  that  the  bulb  when  sealed  was  full  of 


134  LABORATORY  PRACTICE. 

vapour,  break  off  the  tip  of  the  tubulature  under 
water  (recently  boiled  to  drive  out  the  dissolved 
air),  when  the  liquid  should  rush  in  and  complete- 
ly fill  the  interior.  If  any  considerable  volume  of 
residual  air  then  appears  (more  than  five  or  ten 
cubic  centimetres)  the  determination  should  be  re- 
peated, using  more  liquid  and  taking  more  care 
to  seal  the  bulb  at  the  right  time.  If,  as  is  usu- 
ally the  case,  the  material  used  is  combustible, 
the  right  moment  is  easily  caught  by  lighting  the 
jet  of  vapour  as  it  issues  from  the  tubulature 
(after  the  first  violent  rush  has  ceased),  and  watch- 
ing the  flame  as  it  diminishes.  The  moment  the 
flame  disappears  the  tubulature  should  be  sealed. 
In  repeating  the  experiment  it  is  of  course  unne- 
cessary to  alter  the  tare  or  disturb  the  equilibri- 
um if  only  the  tips  broken  off  are  kept  and  re- 
turned to  the  balance  pan  with  the  bulb. 

(1)  The  student  should  now  answer  the  following  ques- 
tions :  (1.)  What  was  the  density  of  the  vapour  which  filled 
the  bulb  when  in  the  kettle  at  the  temperature  and  pressure 
noted  ?  (2.)  According  to  what  laws  does  the  density  of  a 
dry  vapour  vary  when  it  freely  partakes  of  the  temperature 
and  pressure  of  the  surrounding  medium  ?  (3.)  Does  a  va- 
pour confined  over  the  liquid  from  which  it  rises  conform  to 
the  same  laws  ?  (4.)  Why  does  not  the  specific  gravity  of  a 
vapour  vary  with  the  temperature  and  pressure  ?  (5.)  Why 
is  the  molecular  weight  of  a  substance  equal  to  twice  its 
specific  gravity  in  the  aeriform  state  referred  to  hydrogen 


(2)  The  student  should  carefully  review  in  this  connec- 
tion Subdivision  2,  on  Air. 


MOLECULAR  WEIGHT.  135 

(3)  The  above  method  obviously  only  applies  to  such 
liquids  as  boil  below  the  boiling  point  of  water.     With  less 
volatile  liquids  the  bulb  may  be  sunk,  by  means  of  appro 
priate  apparatus,  in  a  bath  of  melted  paraffine  maintained  at 
a  constant  temperature  ;  and  the  specific  gravity  of  the  va- 
pours of  comparatively  fixed  bodies  has  been  formed  by 
using  globes  of  porcelain  heated  in  a  bath  of  boiling  zinc  by 
means  of  very  powerful  furnaces.     Quite  a  different  method 
of  experimenting  *  is  better  adapted  to  such  cases,  but  it  is 
beyond  the  scope  of  this  book. 

(4)  Sharp  results  conforming  to  theory  can  not  be  ex- 
pected by  the  method  here  described  unless  the  materials 
used  are  of  a  very  high  degree  of  purity.     The  alcohol  must 
be  absolute  and  the  ether  or  chlorform  free  from  all  admixt- 
ures.    Obviously  a  process  of  fractional  distillation  takes 
place  in  the  bulb  and  the  impurities  may  be  concentrated  in 
the  residual  vapour  which  is  weighed.     This  objectionable 
feature  of  the  process  is  avoided  in  still  a  third  method,  ap- 
plicable only  to  comparatively  volatile  liquids  in  which  the 
volume  of  vapour  formed   by  a  weighed  amount  of  sub- 
stance is  accurately  measured  under  observed  conditions. 
For  a  description  of  this  process  see  loc.  cit.  in  (3). 


20.  Chemical  Method  of  determining  Mo- 
lecular Weights. 

From  the  chemist's  point  of  view,  the  molecule 
of  a  substance  is  the  ultimate  particle  which  pos- 
sesses the  qualities  of  the  substance.  Molecules 
may  be  divided,  but  if  divided  the  properties  of 
the  substance  are  lost,  new  molecules  are  formed, 
and  new  substances  with  new  properties  appear. 
In  every  chemical  process  the  change  must  take 

*  See  author's  Chemical  Philosophy,  pp.  33-37. 


136  LABORATORY  PRACTICE. 

place  between  molecules ;  that  is,  one  or  more 
molecules  of  one  substance  must  act  upon  or  must 
yield  one  or  more  molecules  of  another  substance, 
Hence  it  must  be  that  in  any  chemical  change  the 
weights  of  the  substances  involved  must  be  in 
the  proportion  of  their  molecular  weights,  or  in 
some  multiple  of  this  proportion.  In  other  words, 
assuming  the  fundamental  conception  of  molecu- 
lar structure,  the  law  of  definite  proportion  is  a 
necessary  deduction  of  the  molecular  theory.  And 
not  only  is  the  law  of  definite  proportions  a  fixed 
principle  of  nature,  but,  moreover,  the  definite 
proportions  shown  by  chemical  analysis  are  found 
to  bear  a  very  simple  numerical  relation  to  the 
molecular  weights  measured  by  the  vapour  densi- 
ties. Thus  the  facts  of  chemistry  furnish  a  very 
remarkable  confirmation  of  the  molecular  theory 
of  physics.  Moreover,  the  methods  of  quantita- 
tive chemical  analysis  give  us  another  method  of 
determining  molecular  weights  ;  for  if  in  any 
chemical  process  we  can  find  the  quantitative  rela- 
tions between  any  two  of  the  substances  con- 
cerned, whether  as  factors  or  products,  the  ratio 
of  these  weights  must  be  the  proportion  of  the 
molecular  weights,  or  else  some  simple  multiple  of 
that  proportion  ;  and  in  most  cases  we  are  able  to 
infer  from  the  chemical  relations  what  the  multi- 
ple is. 

Ex.  72.  Molecular  Weights  of  Potassic  Ohio- 


MOLECULAR   WEIGHT.  137 

rate  and  of  Potassic  Chloride. — Weigh,  in  a  porce- 
lain crucible  (or  better,  a  small  platinum  crucible) 
about  two  grammes  of  dry  potassic  chlorate  (pow- 
dered). Heat  to  fusion,  gradually  increasing  the 
temperature  as  the  oxygen  gas  escapes,  taking 
care  to  avoid  sputtering,  and  finally  heating  to 
low  redness  for  several  minutes.  Weigh  the  resi- 
due. Make  the  proportion,  As  the  weight  of  the 
oxygen  driven  off  is  to  the  weight  of  potassic 
chlorate  taken,  so  is  32,  the  molecular  weight  of 
oxygen  gas  assumed  to  be  known,  to  the  corre- 
sponding weight  of  potnssic  chlorate.  The  weight 
thus  found  is  Known  to  be  two  thirds  of  the  mo- 
lecular weight.  What  is  the  molecular  weight1? 
Make  also  the  proportion,  As  the  weight  of  oxy- 
gen expelled  is  to  the  weight  of  potassic  chloride 
left,  so  is  32,  as  before,  to  a  number  which  is 
known  to  be  two  thirds  of  the  molecular  weight  of 
potassic  chloride.  What  is  the  molecular  weight 
of  potassic  chloride  ? 

Ex.  73.  Molecular  Weight  of  Oxalic  Acid.— 
To  a  small  and  light  glass  flask  fit  a  cork  with  two 
perforations  ;  to  one  of  these  adapt  a  small  chlo- 
ride-of-calcium  tube,  and  to  the  other  a  short  out- 
let tube.  Weigh  in  the  flask  about  one  gramme 
of  crystallized  oxalic  acid,  determining  the  weight 
with  precision.  Dissolve  in  fifty  cubic  centime- 
tres of  water  and  add  ten  cubic  centimetres  of 
strong  sulphuric  acid ;  allow  to  cool,  and  then 


138  LABORATORY  PRACTICE. 

add  one  gramme  of  powdered  pyrolusite.  After 
the  apparatus  is  thus  mounted  take  the  tare. 
Then,  closing  the  outlet  tube  with  a  small  bit  of 
wax,  gently  heat  the  flask  so  long  as  the  evolu- 
tion of  carbonic  dioxide  continues.  Again  allow 
to  cool,  and  when  cold  remove  the  wax  stopper 
and  suck  through  the  chloride-of-calcium  tube  so 
long  as  the  taste  of  carbonic  acid  is  perceptible. 
Lastly,  determine  the  loss  of  weight,  which  is  the 
weight  of  the  carbonic  dioxide  formed  in  this 
somewhat  complex  chemical  process.  Still  the 
principle  holds  that  as  the  weight  of  the  carbonic 
dioxide  formed  is  to  the  weight  of  the  oxalic  acid 
used  so  is  44,  the  known  molecular  weight  of  car- 
bonic dioxide,  to  a  number  which  must  be  a  sim- 
ple multiple  or  submultiple  of  the  molecular 
weight  of  crystallized  oxalic  acid.  In  this  case 
the  result  will  be  one  half  of  the  required  molecu- 
lar weight.  Let  the  student  see  that  in  this 
method  it  is  only  necessary  to  know  the  relative 
weights  of  two  of  the  substances  concerned  in  the 
chemical  process,  which  may  be  very  complex 
and  in  regard  to  which  nothing  else  need  be 
known.  Let  him  also  notice  that  the  chemical 
method  has  the  advantage  over  the  physical 
method  in  that  it  is  applicable  to  substances  which 
are  not  volatile.  It  adopts  the  same  unit  as  the 
physical  method,  and  refers  the  unknown  molecu- 
lar weight  to  a  molecular  weight  previously  deter- 


ATOMS.  139 

mined  and  in  the  first  instance  controlled  by  the 
physical  method  ;  but  in  this  way,  step  by  step,  it 
covers  the  whole  ground.  It  is  far  more  accurate 
than  the  physical  method,  and  practically  the 
physical  method  is  only  used  to  control  the  re- 
sults of  chemical  analysis  —  that  is,  to  show 
whether  the  definite  proportions  observed  are  the 
relations  between  single  molecules  or  between 
multiple  molecules.  Since,  however,  the  chem- 
ical method  involves  the  question  of  multiple 
ratios,  it  has  its  necessary  limitations.  When  the 
substance  under  examination  is  volatile,  the  re- 
sults of  analysis  can,  as  just  said,  be  controlled 
by  a  determination  of  vapour  density.  In  other 
cases  a  study  of  the  chemical  process  itself  gives 
us  the  additional  information  required,  but  in  a 
way  that  can  not  be  made  intelligible  in  this  con- 
nection. Not  unfrequently,  however,  all  these 
means  fail,  and  then  the  chemist  is  forced  to  se- 
lect between  the  possible  multiples  the  value 
which  he  thinks  most  probable. 

21.   Conception  of  Atoms. 

The  ultimate  particles  of  substances,  called 
molecules,  although  far  beyond  the  range  of  visi- 
bility, are  by  no  means  inconceivably  small,  and 
still  less  indivisible ;  and  the  teacher  should  at- 
tempt to  aid  the  student's  imagination  by  giving 


140  LABORATORY  PRACTICE. 

the  estimates  of  physicists  in  regard  to  the  abso- 
lute size  of  these  bodies,  and  showing  that  the 
relations  between  their  dimensions  and  our  ordi- 
nary standards  of  magnitude  are  no  more  extreme 
than  those  we  meet  with  in  astronomy,  in  elec- 
tricity, and  in  other  branches  of  physical  science. 
So  far  are  the  molecules  from  being  indivisible 
that  it  is  perfectly  evident  that  they  must  be  di- 
vided in  almost  every  chemical  change.  For  ex- 
ample, as  we  have  seen,  two  volumes  of  hydrogen 
gas  combined  with  one  volume  of  oxygen  gas  to 
form  two  volumes  of  aqueous  vapour.  Here  it  is 
evident  that  if  equal  gas  volumes  contain  the  same 
number  of  molecules,  it  must  be  that  every  two 
molecules  of  hydrogen  gas  combine  with  one 
molecule  of  oxygen  gas  to  form  two  molecules  of 
water — that  is  to  say,  the  molecule  of  oxygen  is 
divided  between  two  molecules  of  water ;  or, 
again,  every  molecule  of  water  contains  one  half 
as  much  oxygen  as  the  molecule  of  oxygen  gas. 
We  reach  the  same  result  in  the  analysis  of  water. 
If  we  calculate  from  the  results  of  analysis  the 
percentage  composition  of  water,  we  find  that  in 
100  parts  water  contains  of  hydrogen  11 -111  per 
cent,  and  of  oxygen  88*888  per  cent.  Further, 
the  specific  gravity  of  aqueous  vapour  referred  to 
hydrogen  is  9,  and  hence  the  molecular  weight  of 
water  is  18.  Of  this  weight,  88 '888  per  cent,  or  16 
parts,  consist  of  oxygen.  Again,  the  specific 


ATOMIC  WEIGHT. 

gravity  of  oxygen  gas  referred  to  hydrogen  is  16, 
and  therefore  the  molecular  weight  of  oxygen  gas 
is  32.  We  know  then  that — 

One  molecule  of  oxygen  gas  weighs  32  microcriths. 

One  molecule  of  water  weighs  18  microcriths. 

One  molecule  of  water  contains  16  microcriths  of  oxygen. 

Hence,  as  before,  every  molecule  of  water  con- 
tains one  half  as  much  oxygen  as  the  molecule  of 
oxygen  gas.  Obviously  we  can  repeat  this  calcula- 
tion with  every  compound  of  oxygen  in  regard  to 
which  we  know  the  molecular  weight  and  the  per 
cent  of  oxygen  which  the  compound  contains.  If 
now  we  arrange  the  results  in  a  table,  as  below — 

ATOMIC    WEIGHT    OF    OXYGEN. 

Compounds  of                   Observed  Weight  of  Mole-       Weight  of 

Oxygen.                           Sp.  Gr.  cule.  Oxygen  in 

Molecule. 

M.  c.  M.  c. 

Water 9'00  18  16 

Carbonic  oxide 13 '95  28  16 

Nitric  oxide 14*97  30  16 

Alcohol 23'28  46  16 

Ether 37'32  74  16 

Carbonic  dioxide 22 '06  44  32 

Nitric  dioxide 24 '82  46  32 

Sulphurous  dioxide. ...     32 '24  64  32 

Acetic  acid 30'07  60  48 

Sulphuric  trioxide 39 '87  80  48 

Osmic  tetroxide 128 '30  263 '2  64 

Oxygen  gas 15'96  32  32 

it  will  appear  that  in  every  case  the  molecule  of 

an  oxygen  compound  contains  either  sixteen  mi- 
10 


142  LABORATORY  PRACTICE. 

crocriths  of  oxygen  or  a  simple  multiple  of  six- 
teen microcriths.  The  smallest  amount  of  oxygen 
in  any  molecule  is  sixteen  microcriths,  and  this  is 
the  weight  of  what  we  call  an  atom  of  oxygen. 
The  word  "atom"  is  derived  from  a  Greek  word 
meaning  indivisible ;  and  this  smallest  known 
mass  of  oxygen,  weighing  sixteen  microcriths,  has 
never  been  divided.  The  molecule  of  oxygen  gas 
consists  of  two  atoms,  and  of  course  can  be  broken 
in  two.  We  can  reason  in  regard  to  the  com- 
pounds of  every  other  elementary  substance  in 
precisely  the  same  way  and  make  a  similar  table,* 
and  in  each  case  we  shall  find  that  the  weights  of 
a  given  element  in  the  molecules  of  its  several 
compounds  are  simple  multiples  of  a  smallest 
weight,  which  we  take  as  the  weight  of  the  atom 
of  that  element.  In  the  case  of  the  compounds  of 
hydrogen  the  smallest  weight  is  one  microcrith, 
which  is  the  weight  of  the  atom  of  hydrogen — the 
smallest  mass  of  matter  recognized  by  science.  It 
is  not  unreasonable,  therefore,  that  it  should  be 
selected  as  the  unit  of  molecular  and  atomic 
weights  ;  and  we  call  this  unit  by  a  definite  name, 
a  microcrith,  in  order  that  the  student  may  asso- 
ciate with  the  name  a  real  if  not  a  tangible  mass 
of  matter.  The  molecule  of  hydrogen  gas,  like 
the  molecule  of  oxygen  gas,  contains  two  micro- 

*  See  author's  Chemical  Philosophy,  pp.  43-45,  or  New  Chemis- 
try, pp.  141. 


ATOMIC  WEIGHT.  143 

criths,  and  hence  before  we  attained  the  concep- 
tion of  the  hydrogen  atom  we  described  correctly 
the  unit  of  molecular  weight  as  the  half -hydrogen 
molecule.  In  this  way  our  conception  of  atoms 
and  our  general  knowledge  of  atomic  weights 
have  been  reached,  and  in  every  work  on  chemis- 
try the  values  of  the  atomic  weights  adopted  are 
given  in  tables  opposite  to  the  names  of  the  ele- 
ments. (See  table  at  end  of  this  book.)  The 
student  must  seek  to  make  clear  to  his  mind  the 
distinction  between  the  conception  of  the  atom 
and  the  conception  of  the  molecule.  The  ultimate 
particles  which  retain  the  qualities  of  a  substance 
are  molecules,  and  there  are  as  many  kinds  of 
molecules  as  there  are  substances.  Atoms  are  the 
smallest  masses  of  the  chemical  elements  yet 
known,  and  there  are  only  as  many  kinds  of 
atoms  as  of  elements.  To  speak  of  an  atom  of  a 
substance,  especially  of  a  compound  substance,  is 
a  misuse  of  terms.  In  the  case  of  elementary 
substances  we  have  still  to  distinguish  between 
the  molecules  of  the  substances  and  the  atoms  of 
the  element.  There  is  but  one  kind  of  atom  of 
any  element,  but  there  may  be  several  distinct 
elementary  substances.  Thus  in  the  case  of  oxy- 
gen we  have  oxygen  gas,  the  molecules  of  which 
consist  of  two  atoms  of  oxygen,  and  ozone,  a 
wholly  different  substance,  the  molecules  of 
which  consist  of  three  atoms  of  oxygen.  The 


144  LABORATORY  PRACTICE. 

molecules  of  elementary  substances  are  formed  by 
the  aggregation  of  atoms  of  the  same  kind  ;  the 
molecules  of  compound  substances  by  the  aggre- 
gation of  atoms  of  different  kinds.  There  are  a 
few  cases,  like  the  vapours  of  mercury  and  zinc, 
where  the  molecule  consists  of  a  single  atom.  In 
a  chemical  change  the  molecules  of  the  substance 
we  call  the  factors  break  up  into  atoms,  which 
group  themselves  into  new  associations  to  form 
molecules  of  the  products. 

22.  Determination  of  Atomic  Weights. 

The  exact  determination  of  an  atomic  weight 
now  resolves  itself  into  a  simple  question  of 
quantitative  analysis.  If  in  any  process  of  quan- 
titative analysis  we  can  determine  the  weights  of 
two  of  the  elementary  substances  involved,  the 
proportion  between  these  quantities  will  be  either 
the  ratio  of  the  atomic  weights  of  the  two  ele- 
ments, or  else  that  of  some  simple  multiple  of 
these  weights,  the  multiple  in  all  cases  being  pre- 
viously known  from  the  relations  of  the  com- 
pounds of  the  element  as  exhibited  in  such  tables 
as  we  have  described,  or  otherwise. 

Ex.  74.  Atomic  Weight  of  Zinc. — Adapt  to  a 
small  flask  (100  cubic  centimetres)  a  tightly  fitting 
cork  and  exit  tube  leading  to  a  pneumatic  trough. 
Place  in  the  flask  10  cubic  centimetres  of  strong 


ATOMIC  WEIGHT.  145 

hydrochloric  acid  and  20  cubic  centimetres  of 
water.  Clean  scrupulously  a  strip  of  the  purest 
sheet  zinc,  and  accurately  determine  its  weight, 
which  should  be  as  nearly  1*25  gramme  as  pos- 
sible. Use  as  a  receiver  a  glass  flask  of  500  cubic 
centimetres  capacity.  When  all  is  ready,  the 
flask  filled  with  water  standing  inverted  on  the 
shelf  of  the  trough  and  the  mouth  of  the  exit  tube 
under  its  lip,  drop  the  metal  into  the  acid  and 
quickly  cork  the  flask.  This  amount  of  metal 
should  yield  nearly  500  cubic  centimetres  of  hy- 
drogen gas  at  the  ordinary  temperature  of  the 
laboratory.  When  the  apparatus  is  in  equilib- 
rium notice  whether  any  water  has  been  sucked 
back  towards  the  flask.  If  so,  make  the  neces- 
sary allowance  in  measuring  the  volume  of  gas 
formed.  Observe  the  thermometer  and  barome- 
ter. Sinking  now  the  flask  in  the  water  of  the 
trough  until  the  level  of  the  water  is  the  same 
within  and  without  the  glass,  place  the  palm  of 
the  hand  under  the  mouth  and  quickly  invert  the 
flask  and  place  it  on  the  pan  of  the  balance  with 
the  water  it  still  holds.  Take  the  tare,  and,  hav- 
ing exactly  filled  the  flask  with  water,  again 
weigh.  The  difference  of  these  weights — that  is, 
the  weight  in  grammes  of  the  water  required  to 
fill  the  flask,  reduced  for  temperature  if  more  ac- 
curacy is  required — will  give  the  volume  of  the 
gas  collected  at  the  temperature  and  pressure  ob- 


146  LABORATORY  PRACTICE. 

served.  Reduce  the  volume  to  standard  condi- 
tion,  making  allowance  for  the  tension  of  aqueous 
vapour  (Ex.  21  and  Ex.  22).  From  this  volume 
calculate  the  weight  of  hydrogen  formed.  Make 
then  the  proportion,  As  the  weight  of  hydrogen 
is  to  the  weight  of  zinc,  so  is  unity  (the  atomic 
weight  of  hydrogen)  to  a  value  which  we  know, 
from  a  comparison  of  the  zinc  compounds,  to  be 
one  half  of  the  atomic  weight  of  zinc.  Double 
the  value  to  find  the  accepted  atomic  weight. 

In  a  similar  way  the  atomic  weight  of  magne- 
sium may  be  found ;  and  by  dissolving  alumi- 
num in  a  solution  of  caustic  soda  the  atomic 
weight  of  aluminum  may  be  determined  with 
great  accuracy.  As  the  atomic  weights  are  fun- 
damental constants  in  chemical  calculations,  it  is 
essential  that  they  should  be  known  with  the 
greatest  possible  precision  ;  and  hence  a  great 
deal  of  labor  has  been  spent  on  the  analytical 
processes  used  in  determining  their  value.  These 
processes  admit  of  very  different  degrees  of  accu- 
racy. There  are  only  a  very  few  which  in  the 
most  skilful  hands  yield  results  accurate  to  the 
ten-thousandth  part  of  the  quantity  estimated ; 
and  even  the  thousandth  part  is  regarded  as  a 
very  high  degree  of  accuracy  in  chemical  analy- 
sis. Most  processes  do  not  give  results  which  can 
be  relied  upon  much  within  one  per  cent ;  and  in 
many  cases  we  are  forced  to  use  methods  that  are 


ATOMIC  WEIGHT.  147 

far  less  accurate  even  than  this.  In  fixing  the 
precise  value  of  an  atomic  weight  our  choice  is 
usually  limited,  both  as  to  the  material  to  be 
analyzed  and  the  analytical  process  to  be  used,  to 
one  or  two  methods  ;  but  in  almost  all  cases  there 
are  abundant  analyses  of  compounds  of  the  same 
element  which  are  sufficiently  accurate  to  enable 
us  to  interpret  the  results  obtained. 

(5)  Examples  illustrating  the  above  points  may 
be  multiplied  by  the  teacher.  Thus,  the  produc- 
tion of  silver  bromide  from  silver  or  the  reduction 
of  silver  bromide  to  silver  gives  the  means  of  con- 
necting the  atomic  weight  of  bromine  with  that  of 
silver.  (Compare  Ex.  65  and  Ex  68.)  So  also  the 
reduction  of  silver  nitrate  in  a  porcelain  crucible 
on  simple  ignition  will  enable  the  student  to  de 
duce  the  molecular  weight  of  silver  nitrate  from 
the  atomic  weight  of  silver,  and  by  inference  from 
what  will  soon  appear  he  can  thus  determine  also 
the  molecular  weight  of  nitric  acid. 

(1)  It  will  be  observed  that  the  method  of  determining 
an  atomic  weight  is  essentially  the  same  as  the  chemical 
method  of  determining  molecular  weights.  In  each  case  the 
method  is  based  on  the  law  of  definite  proportions,  which 
applies  to  elementary  substances  as  well  as  to  compounds, 
only  in  one  case  the  definite  proportions  are  theoretically 
interpreted  as  the  definite  relative  weights  of  atoms,  and  in 
the  other  as  the  equally  definite  relative  weights  of  mole- 
cules. In  all  instances  what  we  can  determine  with  accuracy 
experimentally  is  a  relative  weight.  What  that  relative 
weight  represents  is  a  question  of  interpretation.  The  ratio 


148  LABORATORY  PRACTICE. 

of  the  two  weights  determined  forms  the  first  two  terms  of 
a  proportion  of  which  the  third  term  is  some  known  molecu- 
lar or  atomic  weight.  We  thus  can  connect  one  molecular 
weight  with  another  or  one  atomic  weight  with  another. 
Moreover,  since  we  refer  both  molecular  and  atomic  weights 
to  the  same  unit,  we  can  connect  a  molecular  weight  with 
an  atomic  weight,  as  is  constantly,  in  fact,  done.  Always 
our  results  are  subject  to  interpretation  so  far  as  regards  the 
question  of  multiple  values  (Ex.  73). 


CHAPTER  IV. 

SYMBOLS   AND   NOMENCLATURE. 

23.  Chemical  Symbols. 

THE  initial  letter,  or  letters,  of  its  Latin  name 
are  used  to  represent  one  atom,  and  therefore  the 
atomic  weight,  of  each  chemical  element.  Thus 
H  stands  for  1  microcrith  of  the  element  hydro- 
gen ;  O,  for  16  microcriths  of  the  element  oxygen ; 
C,  for  12  microcriths  of  the  element  carbon.  Sev- 
eral atoms  are  represented  by  means  of  subnumer- 
als  ;  as  Sa,  which  stands  for  2  X  32  =  64  micro- 
criths of  sulphur ;  C18,  which  stands  for  3  X  35*5  = 
106*5  microcriths  of  chlorine.  Molecules  are  repre- 
sented by  writing  together  the  symbols  of  the 
atoms  of  which  they  consist.  Thus,  HaO  stands 
for  a  molecule  of  water  because  each  molecule  of 
water  is  made  up  of  two  atoms  of  hydrogen 
and  one  atom  of  oxygen ;  HaSO4  stands  for  a 
molecule  of  sulphuric  acid,  consisting  of  two 
atoms  of  hydrogen,  one  atom  of  sulphur,  and 
four  atoms  of  oxygen ;  Oa  stands  for  a  molecule 
of  oxygen  gas,  an  aggregate  of  two  molecules 
of  oxygen ;  while  O8  stands  for  a  molecule  of 


150  LABORATORY  PRACTICE. 

ozone  gas,  which,  although  consisting  solely  of 
oxygen,  has  molecules  made  of  three  atoms  in- 
stead of  two,  and  is  a  wholly  different  substance. 
The  molecular  symbol  represents  the  molecular 
weight,  which  is  the  sum  of  the  weights  of  the 
atoms  of  which  the  molecule  consists.  Thus  the 
molecular  weight  of  sulphuric  acid  is  (2xl)  +  32 
+  (16  X  4)  =  98.  From  a  molecular  symbol  we 
can  always  deduce  the  percentage  composition  of 
the  substance  it  represents.  Thus  it  is  obvious 
that  -^  of  a  molecule  of  sulphuric  acid  consists  of 
hydrogen,  ff  of  sulphur,  and  ff  of  oxygen.  The 
substance,  having  the  same  composition  as  the 
molecule,  must  then  contain  in  one  hundred 

parts — 

Hydrogen 2'04 

Sulphur  32-65 

Oxygen  65'31 

100-00 

On  the  principle  of  Avogadro,  all  molecular 
symbols  represent  the  same  volume  in  the  state  of 
gas ;  thus — 

Hydrogen  Gas.  Ozone.  Carbonic  Dioxide. 

Ha  =  2  m.c.         03  =  48  m.c.         CO2  =  44  m.c. 

Water.  Alcohol. 

HaO  =  18  m.  c.         C2H«O  =  46  m.  c. , 

all  represent  the  same  gas  volumes  compared  un- 
der the  same  conditions  of  temperature  and  press- 
ure. It  follows,  then,  that  the  specific  gravity  of 
a  gas  or  vapour  referred  to  hydrogen  can  be  at 


CHEMICAL  SYMBOLS.  151 

once  deduced  from  the  molecular  symbol  by  divid- 
ing the  molecular  weight  by  2.  The  weight  of  a 
litre  of  hydrogen  when  the  barometer  stands  at 
30  inches  and  the  thermometer  at  0°  is  0*0896 
gramme.  At  273°  under  the  same  pressure  the 
weight  would  be  0*0448  gramme.  At  27°  (a  very 
convenient  standard)*  the  weight  would  be  0*0815 
gramme.  Hence  the  weight  of  a  litre  of  any  gas 
or  vapour,  under  either  condition,  may  also  be 
calculated  from  the  molecular  symbol  by  multi- 
plying the  specific  gravity  obtained  as  above  by 
one  of  these  factors. 

*  If  we  select  300°  absolute  temperature— that  is,  27°  C.— for 
our  standard  temperature  and  30  inches  of  mercury  for  the  stand- 
ard pressure  all  reductions  of  gas  volumes  can  be  made  with  the 
greatest  facility.  A  variation  of  one  degree  from  this  standard 
temperature  corresponds  exactly  to  a  variation  of  one  tenth  of  an 
inch  in  the  barometer,  and  the  effect  of  temperature  can  at  once  be 
eliminated  by  altering  to  a  corresponding  extent  the  height  of  the 
mercury  column  measuring  the  pressure.  Moreover,  as  our  observa- 
tions are  almost  invariably  made  at  temperatures  below  the  stand- 
ard (27°  C.,  or  80-6°  F.),  this  correction  is  usually  additive.  As- 
sume, for  example,  that  the  temperature  is  20°  and  the  observed 
height  of  the  barometer  '30-3  inches,  and  it  is  desired  to  reduce  the 
observed  gas  volume  to  the  assumed  standard.  Were  we  to  raise 
the  temperature  to  27°  it  is  obvious  that  we  should  expand  the  gas  ; 
and  to  bring  it  back  to  its  previous  volume  it  would  be  necessary  to 
increase  the  pressure  by  0*7  inch,  which  corresponds,  as  we  have 
stated,  to  seven  degrees.  It  is  thus  evident  that  at  27°  (the  assumed 
standard)  and  at  31  inches  the  volume  of  gas  would  be  the  same 
as  at  20°  and  30-3  inches.  The  problem  is  then  reduced  to  this : 
Given  a  volume  of  gas  at  31  inches,  to  find  what  would  be  its  volume 
at  30  inches ;  and  this  is  obviously  a  very  simple  case  under  Mari- 
otte's  law.  In  like  manner  all  similar  problems  may  be  solved, 
whether  they  relate  to  the  volume  of  a  given  weight  of  gas  or  to 
the  weight  of  a  given  volume. 


152  LABORATORY   PRACTICE. 

The  specific  gravity  referred  to  air,  or  the  den- 
sity under  other  conditions  of  temperature  and 
pressure,  may  now  be  deduced  by  the  methods 
before  described.  Let  the  student  now  answer,  in 
regard  to  each  of  the  molecular  symbols  last 
given,  this  question :  What  information  does  the 
symbol  give  in  regard  to  the  substance  it  repre- 
sents? He  ought  also  now  to  be  able,  without 
further  assistance,  to  reverse  the  reasoning,  and, 
when  the  percentage  composition  and  vapour 
density  are  given,  to  deduce  the  symbol. 

Assume  that  we  have  given  as  the  result  of 
analysis  that  the  percentage  composition  of  abso- 
lute alcohol  is  as  below,  and  that  the  specific 
gravity  of  the  vapour  referred  to  hydrogen  is  23. 
The  weight  of  a  molecule  of  alcohol  is  then  46, 
and  the  amount  of  each  element  in  a  molecule 
is  that  given  in  the  second  column  of  figures. 
Knowing  now  the  weight  of  the  several  atoms, 
we  easily  find  the  number  of  each  kind  in  one 
molecule — 

Carbon  ....  52'17  24  =  2  x  12  Ci 
Hydrogen  .  .  .  13'05  6=6x1  H6 
Oxygen  ....  3478  16  =  1x16  O 

100-00        46 

By  studying  the  above  scheme  it  will  be  seen 
that  we  shall  reach  the  same  result  if  we  at  once 
divide  the  percentages  by  the  atomic  weights,  and 


CHEMICAL  SYMBOLS.  153 

then  seek  the  simplest  ratio  of  whole  numbers 
corresponding  to  the  results. 
Since — 

5217  : 13-05  :  3478  =  24  :  6  : 16, 

it  must  be  that — 

52-17  13-05  34-78   24  6  16  _  g 
12    1    16  - 12  :  1  : 16  ~ 

So  when  we  do  not  know  the  molecular  weight 
we  can  always  find  the  simplest  ratio  of  whole 
atoms  corresponding  to  the  percentage  composi- 
tion, and  the  true  symbol  of  the  compound  must 
be  either  the  symbol  thus  obtained  or  some  mul- 
tiple of  it ;  and  the  molecular  weight  must  be 
either  the  weight  represented  by  this  symbol  or  a 
multiple  of  it.  In  this  way  we  can  always  find  a 
symbol  corresponding  to  the  percentage  composi- 
tion of  minerals  and  of  similar  inactive  and  non- 
volatile chemical  products,  and  we  accept  the 
symbol  thus  obtained  until  further  investigation 
shows  that  some  multiple  of  it  is  more  correct. 
In  practice  the  result  is  often  indefinite,  because 
the  percentages,  owing  to  errors  of  analysis,  are 
not  exactly  known,  and  we  obtain  a  proportion 
which  is  only  approximately  a  simple  ratio  of 
whole  numbers.  In  this  case  we  select  what  we 
regard  as  the  most  probable  ratio,  and  a  good  deal 
of  judgment  is  necessary  in  interpreting  the  re- 
sults. 


154  LABORATORY   PRACTICE. 

(1)  Given  the  percentage  composition  of  chloroform  as 
follows  :  Carhon,   10  '04  ;  hydrogen,   0'83  ;  chlorine,   89  '13. 
Eequired  the  symbol,  knowing  that  the  specific  gravity  of 
chloroform  vapour  equals  59  '75.     Ans.     CHCls. 

(2)  The  percentage  composition  of  sugar  is  given  —  Ex. 
68(2).     What  is  the  symbol  ?    Ans.     C12Ha!.Oii. 

(3)  Calculate  the  percentage  composition  of  nitro-benzol, 
C«HBNOa.     Ans.     Carbon,  58*53  ;  hydrogen,  4*07  ;  nitrogen, 
11-39  ;  oxygen,  26  '01. 

(4)  What  is  the  weight  of  one  litre  of  alcohol  vapour  at 
273°? 

24.    Chemical  Reactions. 

As  a  chemical  process  consists  in  the  breaking 
up  of  the  molecules  of  the  factors  into  atoms  and 
the  regrouping  of  the  same  atoms  without  loss  to 
form  the  molecules  of  the  products,  it  is  obvious 
that  every  chemical  change  may  be  represented 
by  an  equation,  writing  the  symbols  of  the  mole- 
cules of  the  factors  in  the  first  member  and  the 
symbols  of  the  molecules  of  the  products  in  the 
second  member,  using  figures  like  coefficients  in 
algebra  to  indicate  the  number  of  molecules  in- 
volved in  the  process  and  plus  signs  to  separate 
the  symbols.  Thus  the  chemical  change  in  the 
preparation  of  oxygen  gas  from  potassic  chlorate 
(Ex.  23),  is  represented  by  the  equation— 

Potassic  Chlorate.         Potassic  Chloride.        Oxygen  Gas. 
=        2KC1        +        3O, 


Such  an  expression  is  called  in  chemistry  a  re- 
action.    It  presupposes  a  knowledge  of  the  sym- 


CHEMICAL  REACTIONS.  155 

bols  of  the  products  and  of  the  factors  and  also 
of  the  number  of  molecules  which  concur  in  the 
process.  Again,  the  preparation  of  hydrogen  gas 
from  zinc  and  dilute  sulphuric  acid  (Ex.  25)  is  rep- 
resented thus  : 

Dilute  Solution  of  Hydrogen 

Zinc.  Sulphuric  Acid.  Zinc  Sulphate.  Gas. 

Zn     +     (H2SO4  +  Aq)     =     (ZnSCX  +  Aq)     +     Ha 


Here  Aq  indicates  an  indefinite  amount  of  water 
used  to  dilute  the  acid  or  dissolve  the  salt.  So 
also  the  formation  of  ammonia  (Ex.  45)  is  ex- 
pressed by  this  reaction  : 

Nitric  Oxide.  Hydrogen  Gas.  Water.  Ammonic  Gas. 

2  NO        +        5Ha      =      2HaO      +      2  NH, 

In  like  manner  the  student  should  review  all  the 
experiments  he  has  tried,  and  with  the  aid  of  the 
teacher  learn  to  write  the  reaction  in  every  case. 
He  should  be  required  to  state  the  full  signifi- 
cance of  every  reaction  he  writes  until  he  has  ac- 
quired a  complete  mastery  of  this  symbolical  lan- 
guage. 

If  correctly  written,  a  chemical  reaction  illus- 
trates always  the  first  two,  and,  when  it  involves 
aeriform  factors  or  products,  all  three  of  the  fun- 
damental laws  of  chemistry.  The  law  of  con- 
servation of  mass  is  expressed  by  the  equation 
sign,  the  law  of  definite  proportions  by  weight  is 
indicated  by  the  definite  numerical  values  of  the 
several  terms,  and  the  law  of  definite  proportions 


156  LABORATORY  PRACTICE. 

by  volume  is  seen  in  the  simple  ratios  of  the  co- 
efficients. In  like  manner  the  system  of  chemical 
symbols  involves  most  of  the  principles  which 
have  been  discussed  in  this  course  of  experiments. 
For  example,  we  see  at  once  from  the  reaction 
why  in  determining  the  atomic  weight  of  zinc 
in  Ex.  74  the  observed  value  was  doubled,  since 
what  we  observed  was  the  ratio  H2  :  Zn,  and 
what  we  required  the  ratio  H  :  Zn.  So  also  in 
Ex.  66  (b)  we  see  that  water  must  have  been  formed 
by  the  reactions  in  question.  Indeed,  so  fully  do 
the  symbols  embody  the  general  principles  of  chem- 
istry that  students  are  apt  to  infer  that  these  prin- 
ciples have  been  deduced  from  the  symbols,  just 
as  in  mathematics  similar  general  results  have  been 
discovered  by  the  working  out  and  interpretation 
of  algebraic  formulae  ;  and  when,  as  in  many  text 
books,  the  phenomena  are  subordinated  to  the 
symbolical  expression,  this  false  impression  is  in- 
evitable. Chemistry  is  an  inductive,  not  a  de- 
ductive science,  like  mathematics,  and  chemical 
symbols  differ  essentially  from  mathematical  for- 
mulae. In  mathematics  everything  that  can  be 
legitimately  deduced  from  an  algebraic  equation 
must  be  true  ;  but  this  is  far  from  being  the  case 
with  chemical  reactions.  Chemical  symbols  sim- 
ply stand  for  the  facts  and  theories  they  were  de- 
vised to  express  and  for  nothing  more.  To  secure 
the  peculiar  discipline  of  the  physical  sciences  it 


CHEMICAL  REACTIONS.  157 

is  essential  that  they  should  be  studied  as  they 
have  been  built  up.  The  student  must  begin  by 
observing  the  phenomena,  and  be  led  up  to  the 
general  principles  through  his  own  inferences, 
and  this  is  the  order  which  has  been  followed  in 
preparing  this  book.  To  begin  with  an  abstract 
statement  of  these  principles,  or,  what  amounts  to 
the  same  thing,  to  express  at  once  every  phe- 
nomenon observed  in  symbolical  language  which 
embodies  these  principles,  is  to  invert  the  natural 
order  and  to  abandon  the  inductive  method. 
Nevertheless,  such  are  the  perfection  and  grasp  of 
this  system  of  symbols  that  it  is  of  the  greatest 
value  in  aiding  us  to  realize  relations  and  foresee 
results  which  without  it  we  might  not  have  dis- 
covered. 

(1)  The  direction  above  given,  that  the  student  should 
review  all  the  experiments  heretofore  described  and  write  all 
the  reactions  of  the  processes  described  so  far  as  it  is  possi- 
ble, can  not  be  too  strongly  insisted  upon.  Without  such 
practice  the  student  can  not  be  expected  to  grasp  the  sub- 
ject; but  with  it  many  of  the  relations  before  observed  will 
become  clear  and  the  facts  will  all  appear  in  a  clearer  light. 
To  aid  him  we  give  below  the  more  important  reactions,  the 
figure  prefixed  being  the  number  of  the  experiment  under 
which  the  process  is  described  or  indicated  : 

14.  (a)  (H,,C4H40«+HNaCOs+Aq)= 

(HNa,C4H4O6  +  H2O + Aq) + CO, 
14.  (6)   (BaCla+Na2SO4+Aq)=BaS04 
16.         2H90  =  2Ha  +  09 

23.  2KClO3  =  2KCl+3Oa 

24.  (a)  P4  +  5O2  =  2PaO6 
24.  (c)   C  +  Oa  =  CO, 

11 


158  LABORATORY   PRACTICE. 

25.         Zn + (HaSO4  +  Aq)  =  (ZnS04  +  Aq)  +  Hi 
27.         2Ha  +  Oa  =  2HaO 

32.  S  +  Oa  =  SOa 

(SOa  4-  HaO  +  Aq)  =  (HaS08  +  Aq) 

33.  2SOa  +  Oa  =  2SO, 

(SO3  +  HaO  +  Aq)  =  (HaSO4  +  Aq) 

(BaCla  +  HaSO4  +  Aq)  =  BaaS04  4-  (2HC1  4-  Aq) 

(Pa06  4-  3H3O  +  Aq)  =  (2H3PO*  +  Aq) 

34.  (NaCl  +  H2SO4  +  Aq)  =  (HNaSO4  4-  Aq)  +  2HC1 

35.  2HC1  +  Naa  =  2NaCl  +  H9 

36.  (MnOa  4-  4HC1  +  Aq)  =  (MnCla  +  2HaO  +  Aq)  +  Cl, 
H2  +  Cla  =  2HC1 

39.  (a)  (CaCO3  +  2HC1  +  Aq)  =  (CaCla  +  HaO  +  Aq)  +  CO, 

39.  (&)   (CaOaHa  +  Aq)  +  COa  =  (CaCO,  +  (HaO  +  Aq) 

40.  (a)  COa  +  C  =  2CO 

41.  CaH.O  +  HaSO<  =  (HaS04.HaO)  +  CaH4 

CH4=Marsh  Gas,  CaH4=Ethylene,  CaHa= Acetylene. 
43.  (a)  KNO»  +  H3SO4  =  HKSO4  +  HNO3 
43.  (6)  2HNO3  +  S  =  HaSO*  +  2NO 

43.  (c)  2HNO,  +  5Cu  =  5CuO  +  HaO  +  Na 

44.  (a)  3Cu  +  (8HNOs  +  Aq)  = 

(3CuN30«,  +  4HaO  +  Aq)  +  2NO 
2NO  +  Os  =  2NOa 

44.  (&)  10NO  +  P4  =  2PaO6  +  5Na 

45.  2NO  +  5Ha  =  2H,O  +  2NH3 

48.  2Mg  +  Oa  =  2MgO 
MgO  +  H,O  =  MgOaH, 

MgO  +  (HaS04  +  Aq)  =  (MgS04  +  HaO  +  Aq) 

MgOaHa  +  (HaSO4  +  Aq)  =  (MgSO4  +  2H2O  +  Aq) 

Mg  +  (HaSO4  +  Aq)  =  (MgSO4  +  Aq)  +  Ha 

2Zn  +  O2  =  2ZnO 

ZnO  does  not  unite  directly  with  water. 

ZnO  4-  (H2SO4  +  Aq)  =  (ZnSO4  +  HaO  4-  Aq) 

49.  Naa  4-  (2H2O  4-  Aq)  =  (2NaOH  4-  Aq)  +  Ha 
(NaOH  4-  HC1  4-  Aq)  =  (NaCl  4-  HaO  +  Aq) 
(NaOH  +  HNO3  4-  Aq)  =  (NaN03  4-  HaO  4-  Aq) 
(NaOH  +  COa  4-  Aq)  =  (HNaCO.  4-  Aq) 

50.  (c)  Naa  4-  Oil  =  2NaCl 

2Na,  4-  O>  =  2Na,O 


CHEMICAL  REACTIONS.  159 

Na»O  +  H2O  =  2NaOH 

(Na2O  +  2HC1  +  Aq)  =  (2NaCl  +  H2O  +  Aq) 
(NaOH  +  HC1  +  Aq)  =  (NaCl  +  HaO  +  Aq) 
61.         (2NH3  +  H2SO4  +  Aq)  =  ((NH4)aSO4  +  Aq) 
(NH8  +  HNO3  +  Aq)  =  (NH4NO,  +  Aq) 
(NIL  +  HC1  +  Aq)  =  (NH4C1  +  Aq) 
(2NaOH  +  HaSO4  +  Aq)  =  (NaaSO4  +  2HaO  +Aq) 
(NaOH  +  HNO3  +  Aq)  =  (NaNOa  +  H,O  +  Aq) 
(NaOH  +  HC1  +  Aq)  =  (NaCl  +  HaO  +  Aq) 

53.  (a)  Cu  +  O  =  CuO 

CuO  +  Ha  =  Cu  +  HaO 

54.  (a)  (CuO  +  HaSO4  +  Aq)  =  (CuSO4  +  HaO  +  Aq) 

(CuO  +  2HNOs  +  Aq)  =  (CuN2O«  +  HaO  +  Aq) 
(CuO  +  2HC1  +  Aq)  =  (CuCla  +  HaO  +  Aq) 

57.  (a)  Fe  4-  H2S04  +  Aq)  =  (FeS04  -I-  Aq)  +  Ha 

58.  FeO  =  Ferrous  Oxide 
FeaO3  =  Ferric  Oxide 

FeOSOs  (or  FeSO4)  =  Ferrous  Sulphate 
Fea03.3S03  =  Ferric  Sulphate 

59.  (a)  Fe  +  S  =  FeS 

59.  (6)  FeS  +  (HaSO4  +  Aq)  =  (FeSO4  +  Aq)  +  HaS 
FeO  +  (H2SO4  +  Aq)  =  (FeSO*  +  Aq  +  HaO) 
65.         (AgNOs  +  HC1  +  Aq)  =  AgCl  +  (HNO.  -I-  Aq) 

2AgCl  +  Ha  =  2Ag  +  2HC1 

68.         (AgNOs  +  KBr  +  Aq)  =  (KNO8  +  Aq)  +  AgBr 
73.        (MnOa  +  H3SO4  +  H2CaO4  +  Aq)  = 

(MnSO4  +  2H2O  +  Aq)  +  2CO3 

(2)  To  ensure  a  full  comprehension  of  the  subject  the 
teacher  should  ask  such  questions,  as  the  following  : 

Does  it  appear  from  27  that  when  hydrogen  unites 
with  oxygen  two  volumes  of  the  first  combine  with  one 
volume  of  the  second  to  form  two  volumes  ? 

Does  it  appear  that  when  either  charcoal  or  sulphur  burn 
in  oxygen  gas  the  volume  of  the  product  is  the  same  as  the 
volume  of  the  oxygen  consumed  ? 

In  the  ordinary  process  of  preparing  oxygen  gas,  how 
does  the  residual  salt  differ  in  composition  from  the  salt 
used  ? 

Both  zinc  and  zinc  oxide  when  dissolved  in  dilute  sul- 


160  LABORATORY  PRACTICE. 

phuric  acid  yield  the  same  zinc  sulphate.  Why  is  hydrogen 
gas  evolved  in  the  first  case  and  not  in  the  second  ? 

How  do  you  explain  the  production  of  nitric  acid  from 
nitre— Ex.  43  (a)  ? 

When  iron  dissolves  in  nitric  acid  either  nitric  oxide  or 
nitrogen  gas  is  evolved  (as  in  case  of  copper),  while  from 
dilute  sulphuric  acid  the  same  metal  liberates  hydrogen. 
Why  the  difference  ? 

When  to  a  solution  of  a  silver  coin  we  add  hydrochloric 
acid  all  the  silver  is  precipitated  as  chloride,  but  none  of  the 
copper.  Why  this  selection  ? 

By  the  judicious  use  of  such  questions  the  student  will 
be  led  to  think,  the  deadening  effect  of  mechanical  routine 
will  be  avoided,  and  the  teaching  power  of  the  course  greatly 
increased. 

25.  Stochiometry. 

Since  in  writing  a  chemical  reaction  the  rela- 
tive weights  of  all  the  factors  and  products  are  ne- 
cessarily implied,  it  follows  that  if  the  total 
weight  of  any  one  substance  concerned  in  the 
process  is  given  the  weight  of  every  other  may  be 
calculated.  It  is  only  necessary  to  make  the  pro- 
portion— 

As  the  total  molecular  weight  of  the  given  sub- 
stance is  to  the  total  molecular  weight  of  the  re- 
quired substance,  so  is  the  gross  weight  given  to 
the  gross  weight  required. 

By  total  molecular  weight  is  here  meant  the 
simple  molecular  weight  of  the  substance  multi- 
plied by  the  coefficient  with  which  it  appears  in 
the  reaction.  If  in  such  problems  a  volume  is 


STOCHIOMETRY.  161 

given  this  volume  must  be  reduced  to  weight  by 
the  simple  method  already  described  before  ap- 
plying the  rule  ;  and  when  the  volume  of  a  factor 
or  product  is  sought  the  reverse  reduction  is 
readily  made  after  the  weight  is  known.  The  re- 
lation between  any  two  gas  volumes  is  of  course 
directly  seen  on  inspecting  the  reaction,  and  needs 
no  calculation.  The  student  ought  to  have  a 
great  deal  of  practice  in  stochiometrical  calcula- 
tion. A  very  large  number  of  problems  of  this 
sort  will  be  found  in  the  author's  Chemical  Phi- 
losophy, and  there  are  many  works  wholly  de- 
voted to  the  subject.  The  teacher,  however,  will 
add  interest  to  a  necessarily  dry  subject  if  he 
constructs  problems  of  his  own,  based  on  the  ex- 
periments which  the  student  has  actually  per- 
formed. 

(1)  How  many  grammes  of  common  salt  can  be  made 
from  25  grammes  of  sodium  bicarbonate  ? 

(2)  In  the  process  of  making  nitric  acid,  how  many 
grammes  of  sulphuric  acid  will  be  required  to  every  kilo- 
gramme of  nitre,  assuming  that  the  acid  used  contains  95 
per  cent  of  HaSCh  ?    How  many  cubic  centimetres  would  be 
required  ?    (Such  an  acid  has  the  specific  gravity  T84.) 

(3)  How  many  grammes  of  charcoal  and  how  many  of 
sulphur  will  burn  in  one  litre  of  oxygen  measured  at  stand- 
ard conditions  ?    What  will  be  the  weight  of  one  litre  of 
each  of  the  products  under  the  same  conditions  ? 

(4)  How  many  grammes  of  potassium  chlorate  are  re- 
quired to  yield  four  litres  of  dry  oxygen  (standard  condi- 
tions) ? 

(5)  When  dissolved  in  acid  1*25  gramme  of  zinc  will 
yield  what  volume  of  hydrogen  gas  collected  over  water 


162  LABORATORY  PRACTICE. 

when  temperature  of  room  is  20  and  barometer  stands  at 
750  millimetres  ? 

(6)  In  preparing  ammonia  gas  from  nitric  oxide  and  hy- 
drogen gas  in  what  proportions  by  volume  should  the  last 
two  be  mixed  ? 

(7)  In  preparing  nitric  oxide,  how  many  grammes  of 
copper  will  be  required  to  each  litre  of  gas  if  the  product 
is  only  nitric  oxide  ?    How  many  if  the  product  is  wholly 
nitrogen  gas  ? 

(8)  A  cubic  decimetre  of  marble  contains  how  many 
times  its  own  volume  of  carbonic-dioxide  gas  ?     Specific 
gravity  of  marble,  2*75. 

(9)  The  reactions  involved  in  these  problems  are  all 
given  in  the  list  under  the  preceding  division     The  student 
should  not  limit  his  study  to  the  few  problems  here  given  as 
examples,  and  it  must  be  borne  in  mind  that  the  same  prin- 
ciples apply  to  any  chemical  process,  however  complex  or 
however  so  many  simple  reactions  it  may  involve,  provided 
only  that  the  whole  material  from  one  reaction  passes  for- 
ward to  the  next. 


26.  Nomenclature. 

Before  the  present  century  the  names  given  to 
chemical  products  were  almost  wholly  arbitrary, 
and  a  few  of  these,  like  oil  of  vitriol,  blue  vitriol, 
sugar  of  lead,  calomel,  and  Epsom  salts,  still  re- 
main in  common  use.  In  1787  a  systematic  no- 
menclature was  devised  by  a  committee  of  the 
French  Academy  of  Sciences,  under  the  lead  of 
Lavoisier,  in  which  the  name  of  a  substance  was 
made  to  indicate  its  composition,  and  at  the  time 
of  its  adoption  and  for  more  than  fifty  years  after- 
wards it  was  probably  the  most  perfect  nomen- 


NOMENCLATURE.  163 

clature  which  any  science  ever  enjoyed.  It  was 
based,  however,  on  the  dualistic  theory  of  La- 
voisier, and  when  the  science  outgrew  the  theory 
the  old  names  lost  much  of  their  significance  and 
appropriateness.  Nevertheless,  the  main  features 
of  the  Lavoisierian  nomenclature  are  still  pre- 
served, although  with  some  variations  of  usage  as 
to  details  and  the  introduction  of  many  arbitrary 
names,  like  carbinol,  phenol,  pinakone,  de- 
manded by  the  necessities  of  a  rapidly  expanding 
science.  The  old  nomenclature  is  so  oat  of  har- 
mony with  our  modern  conceptions  that  it  would 
be  impossible  to  explain  its  full  significance  with- 
out entering  into  details  which  would  be  out  of 
place  in  an  elementary  course.  A  few  of  the  rules 
should  be  stated  by  the  teacher  and  the  use  of 
the  ordinary  terminations  and  prefixes  so  far  ex- 
plained as  to  render  the  usually  occurring  names 
intelligible.  All  that  is  required  may  be  found  in 
any  elementary  text  book  on  chemistry. 


CHAPTER  V. 

MOLECULAK   STKUCTUBE. 

27.  Quantivalence. 

OF  all  chemical  reactions  by  far  the  most  com- 
mon is  a  class  in  which,  judging  from  the  prod- 
ucts, the  only  change  that  takes  place  is  an  inter- 
change of  atoms  or  groups  of  atoms  between  two 
sets  of  molecules,  leaving  all  relations  otherwise 
the  same  as  before.  Such  reactions  are  described 
as  metathetical,  and  the  process  is  termed  meta- 
thesis. Our  chemical  symbols  here  come  to  our 
aid  by  enabling  us  to  form  a  clear  idea  of  what  is 
meant  by  these  terms.  Thus  in  the  reaction  of  a 
solution  of  silver  nitrate  or  a  solution  of  potas- 
sium bromide  (Ex.  68)— 

(AgNOs  -I-  KBr  +  Aq)  =  (KNO3  +  Aq)  +  AgBr 

it  is  obvious  that  Ag  changes  place  with  the  K. 
So  also  in  the  reaction  by  which  hydrogen  gas  is 
made  from  zinc  and  dilute  sulphuric  acid— 

(H2SO*  +  Aq)  +  Zn  =  (ZnSO*  +  Aq)  +  H, 
it  is  equally  obvious  that  Zn  changes  place  with 


METATHESIS.  165 

H,.  Again  in  the  ordinary  test  for  sulphuric 
acid — 

(H,SO4  +  Bad,  +  Aq)  =  BaSO4  +  (2  HC1  +  Aq) 

it  is  evident  that  Ba  has  changed  place  with  H,. 
The  following  experiment  will  further  illustrate 
this  point : 

Ex.  75.  Metathesis. — Pour  ten  cubic  centime- 
tres of  water  into  a  test  tube  and  dissolve  in  it 
one  gramme  of  silver  nitrate.  Immerse  in  the 
liquid  a  small  strip  of  pure  copper,  whose  weight 
must  be  accurately  determined  and  must  not  ex- 
ceed eighteen  centigrammes.  After  the  silver  has 
separated,  wash  the  powder  on  to  a  small  filter 
and  continue  to  pour  water  on  to  the  filter  until  it 
runs  through  tasteless.  Dry  the  filter  on  the  tun- 
nel. Remove  when  dry,  and  after  carefully  wrap- 
ping the  loose  paper  round  the  silver  powder  place 
the  ball  in  a  tared  porcelain  crucible  and  slowly 
heat  to  redness  until  the  paper  has  been  burned, 
when  the  silver  will  appear  bright.  When  cold 
again  weigh  the  crucible  and  find  the  weight  of 
the  silver,  which  should  be  to  the  weight  of  the 
copper  approximately  as  Ag2  =  216  :  Cu  =  63 '6. 

Immerse  now  in  the  blue  liquid  decanted  from 
the  silver  a  strip  of  zinc.  The  copper  which  had 
passed  into  solution  in  the  previous  process  will 
now  be  precipitated.  The  reactions  may  be  writ- 
ten: 


166          LABORATORY  PRACTICE. 

Cu  +  (2  AgNO8  +  Aq)  =  Aga  +  (Cu(NO3)2  4-  Aq). 
Zn  +  (Cu(N08)3  +  Aq)  =  Cu  +  (Zn(NO.)«  +  Aq), 

Obviously,  in  the  first,  one  atom  of  copper  re- 
places two  atoms  of  silver,  and  in  the  second  one 
atom  of  zinc  replaces  one  atom  of  copper ;  and  in 
studying  these  reactions  it  must  be  remembered 
that  the  symbols  correctly  represent  the  atomic 
relations.  Now,  by  studying  in  a  similar  way  a 
very  large  number  of  metathetical  reactions,  it  ap- 
pears that  the  atoms  of  hydrogen,  lithium,  so- 
dium, potassium,  caesium,  rubidium,  silver,  thali- 
um,  chlorine,  bromine,  and  iodine  are  alike  in 
this,  that  while  among  themselves  they  can  be  ex- 
changed atom  for  atom,  they  are  replaced  by  all 
other  atoms  in  groups  of  two,  three,  four,  or  more. 
The  atoms  enumerated  appear  to  have  the  small- 
est exchangeable  value  and  are  said  to  be  univa- 
lent,  while  atoms  which  will  fill  the  place  of  two 
univalent  atoms  are  said  to  be  bivalent,  those  that 
can  fill  the  place  of  three  trivalent,  those  that  can 
fill  the  place  of  four  quadrivalent,  etc.  So  also 
the  terms  quanti valence  and  multi valence.  The 
facts  here  stated  suggest  at  once  the  conception  of 
molecular  structure,  for  it  would  seem  as  if  the 
parts  of  a  molecule  must  be  bound  together  in 
definite  relative  positions  in  order  to  render  such 
substitutions  possible.  This  conception  of  struct- 
ure is  greatly  widened  and  strengthened  when  we 
compare  together  the  symbols  of  molecules  formed 


ATOM-FIXING   POWER.  167 

by  the  union  of  atoms  having  different  degrees  of 
quantivalence,  as  shown  in  metathetical  reactions 
similar  to  those  just  described,  for  it  appears  that 
the  combining  power  corresponds  exactly  to  the 
replacing  power.  In  making  such  comparisons  it 
must  constantly  be  borne  in  mind  that  the  symbol 
represents  in  every  case  our  knowledge  of  the 
composition  and  relations  of  the  substance,  and 
that  the  symbolical  language  thus  enables  us  to 
bring  before  the  mind  in  one  view  the  results  of 
long-continued  and  laborious  investigation.  We 
give  below  the  symbols  of  four  well-known  and 
typical  compounds  of  hydrogen  : 

Hydrochloric  Acid.         Water.  Ammonia  Gas.         Marsh  Gas. 

HC1  H20  H3N 


In  these  compounds  the  atoms  Cl,  O,  N,  and 
C  are  united  with  the  same  number  of  univalent 
atoms,  which,  under  other  circumstances,  they 
might  replace.  So,  also,  we  have— 


Common  Salt.       Baric  Chloride.  WorMe          Zirconic  Chloride. 

NaCl  Bad,  SbCls  ZrCl* 

Compare  now  with  these  the  corresponding  ox- 
ides — 

Na,O  BaO  Sb,O,  ZrO,, 

and  it  will  be  seen  how  we  are  led  to  the  conclu- 
sion that  in  these  molecules  the  atoms  of  lower 
quantivalence  are  united  to  those  of  higher  quan- 


H 

H 

i 

i 

H-N-H 

H-C-H 

i 

H 

Cl 

01 

i 

i 

Cl-Sb-Cl 

Cl-Zr-Cl 

i 

01 

168  LABORATORY  PRACTICE. 

tivalence,  which  are,  as  it  were,  the  nucleus  of  the 
molecule,  and  serve  like  a  clamp  to  bind  the  parts 
together.  In  order  to  express  this  we  often  write 
the  symbols  as  below,  using  dashes  to  indicate 
what  we  call  the  bonds.  Symbols  thus  written 
are  said  to  be  graphic,  while  as  before  written 
they  are  not  inappropriately  spoken  of  as  em- 
pirical. 

H-C1          H-O-H 


Na-Cl         Cl-Ba-Cl 


Na-O-Na          Ba=O        O  =  Sb-O-Sb  =  O        O  =  Zr=O 

This  is  but  a  very  short  step  in  our  reasoning,  and 
yet  it  opens  to  view  at  once  a  very  definite  struct- 
ure. Notice  that  the  only  inference  we  have 
drawn  is  that  the  atoms,  instead  of  being  indis- 
criminately piled  together,  are  united  each  sepa- 
rately to  the  multivalent  atom  or  atoms  of  the 
group.  This  inference  granted,  we  can  take  an- 
other step. 

When  sodium  acts  on  water  we  have  a  simple 
metathesis — 

(2H-O-H  +  Aq)  +  Naa  =  (2Na-O-H  +  Aq)  +  Ha, 

and  the  product  Na-O-H  must  have  the  same 
structure  as  H-O-H,  containing  only  Na  in  place 
of  one  of  the  hydrogen  atoms.  In  a  similar  way, 


GRAPHIC  SYMBOLS.  1G9 

when  magnesium  acts  on  water  it  must  be 
that— 

1:8:1  +  Mg  =  Mg:8:i  +  a. 

And  here  it  is  evident  that  the  multivalent  atom 
Mg  binds  together  two  molecules  of  water,  and 
thus  we  can  conceive  how  complex  molecules  may 
be  built  up,  and  we  also  see  that  the  replacing 
and  combining  powers  are  merely  different  mani- 
festations of  the  same  atribute  of  the  atoms  which 
we  express  by  the  term  " bonds";  and  hence 
the  reason  that  the  two  powers  necessarily  cor- 
respond. 

The  products  whose  formation  and  structure 
we  have  studied  are  two  members  of  a  very  large 
class  of  substances,  all  having  similar  relations, 
and  which  must  have  a  similar  structure.  The 
symbols  of  four  other  members  of  the  same  class 

are  given  on  the  next  line : 

-O-H 

K-O-H     Calg'll     Zrlgll     58*3:1 

-O-H  -O-H 

-O-H 

Bodies  of  this  class  are  called  hydrates,  and  the 
chief  feature  in  their  structure  is  that  they  con- 
tain atoms  of  hydrogen  united  with  a  multivalent 
nucleus  through  atoms  of  oxygen;  and,  further, 
the  one  chemical  relation  which  marks  all  such 


1YO  LABORATORY  PRACTICE. 

substances  is  that  the  atoms  of  hydrogen  so 
united  are  easily  replaced  by  simple  metatheti- 
cal  reactions. 

Ex.  76.  Replacement  of  Hydrogen  in  Sodic 
Hydrate. — In  a  small  flask  (50  cubic  centimetres) 
dissolve  5  grammes  of  caustic  soda  in  10  cubic 
centimetres  of  water ;  add  a  strip  of  aluminum 
weighing  less  than  three  decigrammes ;  connect 
with  pneumatic  trough,  and  collect  the  hydrogen 
gas  evolved.  Measure  the  gas  volume,  observing 
thermometer  and  barometer ;  correct  for  tension 
of  aqueous  vapor  and  calculate  the  weight  of 
hydrogen  obtained.  Compare  this  weight  with 
the  weight  of  aluminum,  and  estimate  how  many 
atoms  of  hydrogen  must  have  been  replaced  by 
each  double  atom  of  aluminum  (54  microcriths) 
indicated  by  the  barred  symbol. 

If  now,  in  this  connection,  the  student  will 
review  the  familiar  experiment  by  which  hydro- 
gen gas  is  usually  made,  he  will  see  that  this  pro- 
cess is  also  a  metathetical  reaction  in  which  hy- 
drogen atoms  have  been  replaced  by  atoms  of 
zinc ;  and  if  we  take  such  reactions  as  an  indica- 
tion of  the  type  of  structure  we  have  assigned  to 
the  hydrates,  it  will  appear  on  further  study  that 
most  of  the  active  agents  of  chemistry,  including 
both  acids  and  alkalies,  must  be  classed  in  the 
same  group.  To  class  acids  and  alkalies  in  the 
same  group  of  compounds  seems  at  first  sight 


MOLECULAR  STRUCTURE. 

very  anomalous,  for  in  most  respects  their  prop- 
erties are  the  direct  opposites  each  of  the  other. 
Nevertheless,  not  only  do  these  bodies  resemble 
each  other  in  the  one  essential  relation  of  a  hy- 
drate, but  also  they  may  be  grouped  in  series 
varying  so  gradually  from  strong  alkalies  at  one 
end  to  strong  acids  at  the  other  that  no  natural 
dividing  line  can  be  found.  Moreover,  the  dis- 
tinction between  an  acid  and  a  base  is  of  a  rela- 
tive, and  not  of  an  absolute  character,  as  is  shown 
by  the  fact  that  in  such  series  as  have  been  men- 
tioned a  given  member  may  act  as  an  acid  towards 
the  members  at  one  end  of  the  list,  and  as  a  base 
towards  those  at  the  other  end.  If  now  we  study 
the  symbols  of  the  ordinary  acids,  and  arrange 
the  hydrogen  atoms  after  the  pattern  of  a  hy- 
drate, we  shall  get  such  a  result  as  this : 

Nitric  Acid.  Sulphuric  Acid.  Phosphoric  Acid. 

H-O-NO,  §~2~SO,  H-O-PO 

H-0X 

And  it  will  be  seen  that  though  here,  as  before, 
the  hydrogen  atoms  are  united  through  oxygen 
atoms  to  a  multivalent  atom  which  serves  as  the 
nucleus  of  the  molecule,  this  atom  of  high  quan- 
tivalence  acting  as  an  atomic  clamp  is  one  of  a 
group  of  atoms.  Such  groups  are  termed  in 
chemistry  compound  radicals,  and  the  chain  H-O- 
is  called  hydroxyl.  If  the  reasoning  has  been  fol- 


LABORATORY  PRACTICE. 

lowed  thus  far  the  student  will  be  prepared  to 
admit  that  the  relations  of  acids  and  bases  which 
play  such  an  important  part  in  the  science  of  chem- 
istry depend  upon  molecular  structure.  But  why 
the  opposition  between  these  bodies  ?  As  otherwise 
they  have  the  same  structure,  it  is  evident  that 
the  antagonism  must  be  connected  with  the  atoms 
or  compound  radical  with  which  the  hydroxyl 
groups  are  associated  ;  and  when  we  compare  the 
nuclei  of  the  respective  molecules  we  find  a  very 
manifest  difference  between  the  nucleus  of  a 
marked  alkali  and  the  nucleus  of  a  pronounced 
acid.  In  the  alkali  the  nucleus  is  a  metallic  atom, 
like  sodium  or  potassium  ;  in  the  acid  it  is  a  non- 
metallic  atom,  or  a  group  of  such  atoms,  like 
nitrogen  or  sulphur.  Moreover,  we  find  that 
while  it  is  very  easy  to  replace  the  hydrogen 
atoms  either  of  an  alkali  or  of  an  acid  by  atoms 
unlike  the  nucleus,  it  is  difficult  to  replace  them 
by  atoms  similar  to  the  nucleus.  Thus  it  is  diffi- 
cult to  change  Na-O-H  to  Na-O-Na,  but  very 
easy  to  change  it  to  Na-O-Cl  or  Na-O-NO9.  The 
phenomena  point  to  a  polarity  in  the  molecule 
similar  to  that  of  a  magnet,  and  only  by  such 
analogies  can  we  explain  them. 

The  objects  of  this  discussion  have  been,  in  the 
first  place,  to  give  an  idea  of  the  mode  of  reason- 
ing by  which  our  knowledge  of  molecular  struct- 
ure is  reached  ;  and,  in  the  second  place,  to  ex- 


MOLECULAR  STRUCTURE.  173 

plain  the  distinction  between  acids,  alkalies,  and 
salts.  This  distinction,  however  much  it  may  be 
described,  can  not  be  made  clear  except  through 
the  molecular  structure  on  which  it  is  supposed  to 
depend.  Towards  the  first  object  we  have  been 
able  to  advance  only  a  very  few  steps,  but  far 
enough  to  point  out  the  way  by  which  the  com- 
plex structure  of  organic  compounds  has  been  un- 
ravelled and  the  whole  subject  of  organic  chemis- 
try developed.  As  towards  the  second  object, 
we  hope  we  have  been  able  to  make  clear  that 
acids  and  alkalies  have  in  their  larger  relations 
similar  qualities  and  a  similar  structure,  and  that 
they  differ  in  the  character  of  the  nuclei  to  which 
the  hydroxyl  groups  are  united  ;  and,  further,  that 
when  the  hydrogen  atoms  thus  united  are  replaced, 
in  compounds  of  either  class,  by  atoms  opposite 
in  qualities  to  the  atoms  of  the  nuclei,  the  prod- 
ucts are  salts — so  called  because  for  the  most  part 
they  are  bodies  that  can  be  readily  crystallized. 
In  this  connection  there  are  one  or  two  other 
points  to  be  noticed  before  leaving  the  subject. 

As  has  been  shown,  an  acid  most  readily  com- 
bines with  an  alkali  to  form  a  salt ;  and  the  reason 
seems  to  be  that  the  polarity  of  the  molecules  tends 
to  bring  atoms  of  opposite  characters  to  the  two 
ends,  and  determines  a  metathesis  thus— 

Na-O-H  +  H-O-NOa  =  Na-O-NO9  +  H-O-H 
Ca=Oi  =  Ha  +  H,  =  Oa  =  SOa  =  Ca  =  Oa  =  SOa  +  2H-O-H 
12 


174:  LABORATORY  PRACTICE. 

The  number  of  hydroxyl  groups  in  a  hydrate, 
whether  acid  or  alkali,  measures  what  we  call  its 
atomicity.  Thus  sodic  hydrate  is  monatomic,  and 
aluminic  hydrate  hexatomic.  In  an  acid  the 
atomicity  is  frequently  called  basicity,  and  just  as 
the  metallic  atoms  associated  with  the  acid  nuclei 
in  a  salt  are  often  spoken  of  as  basic  radicals,  so 
the  metallic  hydrates  used  for  neutralizing  the 
acid,  as  in  the  above  reaction,  are  often  termed 
bases,  or  the  base  of  the  salt,*  Thus  nitric  acid  is 
monobasic,  sulphuric  acid  is  dibasic,  and  phos- 
phoric acid  is  tribasic.  Hence,  while  nitric  acid 
will  form  only  one  salt  with  sodium,  sulphuric 
acid  will  form  two,  and  phosphoric  acid  will  form 
three.  This  point  is  illustrated  by  the  following 
symbols — 

Na-O-NO, 


Na:8>so' 

H  -0N 
H  -O^PO 
Na-0x 

Na-Oxqn 
Na-0/b0a 

H  -Ox 
Na-O^PO 
Na-0x 

Na-O 
Na-C 
Na-O 

And  the  fact  that  the  several  salts  which  the  sym- 
bols show  to  be  possible  can  be  prepared  is  in 
harmony  with  the  molecular  structure  we  have 
described. 

The  chief  features  of  the  type  of  molecular 

*  The  term  "  base  "  is  used  in  a  broader  sense  than  the  word 
"  alkali,"  and  is  applied  to  any  hydrate  which  will  unite  with  an 
acid. 


MOLECULAR  STRUCTURE.          175 

structure  we  have  unfolded  are  very  strikingly 
illustrated  by  the  formation  and  decomposition  of 
ammonium  nitrate.  The  molecules  of  ammonia  gas 
must  have  (as  we  have  seen)  the  simple  structure- 
EC 

H-N-H; 

but  when  the  gas  dissolves  in  water  the  solution 
acquires  properties  so  closely  resembling  those  of 
a  solution  of  sodium  hydrate  that  we  naturally 
conclude  that  new  molecules  have  been  formed  by 
combination  with  water  having  a  structure  simi- 
lar to  Na  -  O  -  H  ;  that  is— 

H3N  +  H-O-H  =  H4N-O-H. 

When  next  we  neutralize  ammonium  hydrate 
with  nitric  acid,  as  in  Ex.  46  (a),  the  reaction  must 
be  like  the  reaction  of  the  same  acid  on  sodium 
nitrate  given  above,  or— 

H4N-O-H  +  H-O-NO,  =  H4N-O-N02  +  H-O-H. 

Here,  then,  if  our  reasoning  is  correct,  we  have 
the  molecule  of  a  salt  having  as  a  nucleus  N  -  O  -  N, 
but  with  hydrogen  atoms  attached  to  the  nitro- 
gen atom  which  forms  one  pole  of  the  molecule, 
and  oxygen  atoms  united  to  the  nitrogen  atom 
which  forms  the  opposite  pole  ;  and  that  this  is 
the  true  structure  is  indicated  by  the  fact  that 
when  we  simply  heat  the  salt,  the  oxygen  and 
hydrogen  atoms,  thus  for  a  time  kept  apart,  rush 


176  LABORATORY  PRACTICE. 

into  combination,  and  form  molecules  of  water, 
leaving  the  nuclei  free  to  become  the  molecules  of 
a  well-known  substance  called  nitrous-oxide  gas. 

Ex.  77.  Preparation  of  Nitrous  Oxide. — Con- 
nect a  small  flask  (fifty  cubic  centimetres)  by 
means  of  perforated  corks  and  glass  tubes,  first 
with  a  test  tube  and  second  with  a  pneumatic 
trough.  Place  in  the  flask  twenty-five  grammes 
of  ammonium  nitrate,  and  mount  the  apparatus  so 
that  while  the  flask  is  held  by  a  retort  holder  the 
test  tube  may  stand  in  a  beaker  of  water  and  the 
exit  tube  may  open  under  the  mouth  of  a  glass  jar 
standing  full  of  water  and  inverted  on  the  shelf  of 
the  trough.  Cautiously  heat  the  salt  until  it  melts, 
and  then  press  the  heat  until  decomposition  en- 
sues. Water  will  distil  over  and  collect  in  the 
test  tube,  while  nitrous  oxide  will  bubble  up  and 
displace  the  water  in  the  jar.  When  the  jar  is 
filled,  seal  it  and  preserve  the  gas  for  comparison. 

Ex.  78.  Composition  of  Nitrous  Oxide. — Into 
a  jar  of  nitric  oxide,  prepared  as  in  Ex.  44  (a), 
cautiously  pour  100  cubic  centimetres  of  a  con- 
centrated solution  of  green  vitriol  acidified  with 
hydrochloric  acid,  seal  the  jar,  and  shake  the  solu- 
tion with  the  gas  so  long  as  absorption  continues. 
Open  now  the  mouth  of  the  jar  under  water,  and, 
after  comparing  the  residual  gas  volume  with  the 
original  volume  of  the  nitric  oxide,  identify  the 
product  as  the  same  substance  which  was  formed 


MOLECULAR  STRUCTURE.         177 

in  the  last  experiment.  Knowing  that  the  symbol 
of  nitric  oxide  is  NO,  and  that  the  effect  of  the 
green  vitriol  is  to  withdraw  oxygen,  what  infer- 
ence can  you  make  in  regard  to  the  composition  of 
nitrous  oxide  and  as  to  the  nature  of  the  molecu- 
lar change  which  has  taken  place  in  this  experi- 
ment. 

(1)  The  molecular  structure  of  ammonium  nitrate  thus 
developed  may  serve  to  give  some  conception  of  the  condi- 
tions to  which  modern  explosives,  like  nitro-glycerin  and 
gun  cotton,  owe  their  remarkable  relations.  In  the  mole- 
cules of  these  explosives  it  is  supposed  that  three  or  more  of 
the  nuclei  N  -  O  -  N  are  bound  together  by  groups  of  hydrogen 
and  carbon  atoms  at  one  end  of  a  multiple  chain,  while  oxy- 
gen atoms,  in  sufficient  numbers  to  unite  with  all  the  car- 
bon and  hydrogen,  are  attached  at  the  other  end.  By  this 
structure  the  intensely  powerful  affinities  between  the  great 
fire  element  and  the  combustibles  are  for  a  time  held  in 
abeyance.  But  when  the  equilibrium  is  disturbed  the  atoms 
thus  held  apart  rush  together  and  a  great  volume  of  aeriform 
products  are  suddenly  developed  whose  enormous  expansive 
force  produces  the  destructive  effects  so  well  known. 


CHAPTER  VI. 

THERMAL   RELATIONS. 

28.  Heat  of  Chemical  Action. 

EVEN  the  most  elementary  course,  proposing 
to  treat  only  of  the  fundamental  principles  of 
chemistry,  would  be  incomplete  without  some  dis- 
cussion of  the  thermal  relations  of  chemical 
changes,  and  the  most  striking  chemical  experi- 
ments will  have  to  the  student  no  more  meaning 
than  fireworks  if  they  remain  in  his  mind  as  mere- 
ly brilliant  phenomena.  After  what  must  be  here 
assumed  to  have  been  already  learned,  the  student 
should  be  prepared  to  understand  the  treatment 
of  this  subject  in  the  last  chapters  of  the  author's 
New  Chemistry,  and  more  fully  in  the  chapter  on 
the  "Thermal  Eelations  of  Atoms"  in  his  Chemi- 
cal Philosophy.  As  an  introduction  to  the  sub- 
ject, the  student  should  try  the  following  experi- 
ments, using  /for  the  purpose  the  calorimeter, 
already  fully  described  (Ex.  8),  but  when  corro- 
sive liquids  are  used  substituting  for  the  inner 
brass  vessel  as  thin  a  beaker  glass  as  can  be  had, 
of  about  the  same  capacity,  and  filling  the  space 


HEAT  OF  SOLUTION.  179 

between  the  glass  and  the  sides  of  the  chamber 
with  layers  of  wool  wadding,  caught  together  so 
that  the  beaker  can  readily  be  removed  and  re- 
placed. For  accurate  experiments  a  dish  made 
of  thin  platinum  plate  is  always  to  be  preferred. 
In  using  a  beaker  the  heat  absorbed  by  the 
glass  becomes  a  quantity  of  importance.  The 
glass  must  therefore  be  weighed ;  and  the  weight 
of  the  glass  multiplied  by  the  specific  heat  of 
glass  (Ex.  8  (4))  gives  a  value  which  is  called 
the  thermal  water  equivalent,  and  this  in  every 
experiment  is  to  be  added  to  the  weight  of  the 
water. 

Ex.  79.  Heat  of  Hydration  and  of  Solution. 
— Place  in  the  calorimeter  about  300  grammes  of 
water,  weighing  the  amount  accurately  to  a 
gramme.  Prepare  and  pulverize  35  grammes  of 
anhydrous  sodic  sulphate  by  driving  off  the  water 
from  the  crystallized  salt  (Glauber's  salts).  Keep 
the  salt  between  watch  glasses  over  the  beaker  of 
water  and  under  the  cover  of  the  calorimeter  until 
a  perfect  equilibrium  of  temperature  is  reached. 
Then  stir  the  salt  into  the  water  with  a  glass  rod, 
and  observe  the  rise  of  temperature.  Calculate 
the  number  of  units  of  heat  evolved  by  142 
grammes  of  the  salt.*  This  number  is  the  molecu- 

*  In  making  this  and  similar  calculations  it  must  be  remem- 
bered that  is  the  whole  mass  of  the  solution,  and  not  the  water 
merely,  whose  temperature  is  changed.  To  obtain  strictly  accurate 


180  LABORATORY  PRACTICE. 

lar  weight  in  gramme  units  ;  and  it  is  convenient 
to  state  results  on  this  basis,  as  we  can  then 
carry  our  calculations  through  successive  reac- 
tions without  constant  reduction.  Repeat  now 
the  same  experiment,  but  use  79  '3  grammes  Glau- 
ber's salts  (Na2SO* .  10HaO)  in  fine  crystals.  Cal- 
culate as  before  for  one  molecule  in  grammes,  and 
compare  the  two  results.  How  much  heat  is  lib- 
erated in  the  union  of  Na3SO4  with  10HaO  ? 

Ex.  80.  Heat  of  Neutralization. — Weigh  in  a 
glass-stoppered  vial  about  twenty -five  grammes  of 
strong  sulphuric  acid,  the  specific  gravity  of  which 
has  been  previously  accurately  ascertained.  Mix 
this  acid  with  about  250  cubic  centimetres  of  water 
and  give  time  to  cool  before  pouring  into  the  cal- 
orimeter. Finding  from  the  tables  the  amount  of 
HaS04  thus  taken,  calculate  the  amount  of  Na-O- 
H  required  to  neutralize  the  acid  and  weigh  out 
about  one  fifth  more  than  the  calculated  amount 
in  order  to  insure  an  excess.  Dissolve  the  alkali 
also  in  about  250  cubic  centimetres  of  water  and 
place  the  vessels  holding  the  two  solutions  under 
cover  in  a  protected  place  at  the  side  of  the  cal- 

results,  we  should  know  the  specific  heat  of  the  solution ;  but  for  all 
practical  purposes  it  is  sufficiently  accurate  to  count  one  cubic  cen- 
timeter of  the  solution,  measured  at  4°  C.,  as  the  thermal  equiva- 
lent of  one  gramme  of  water.  The  simplest  way  is,  after  the  close 
of  the  experiment,  to  weigh  the  solution  and  determine  its  specific 
gravity  with  a  delicate  spindle  hydrometer  graduated  at  4°  C.  Then 
the  weight,  divided  by  the  specific  gravity,  is  the  thermal  water 
equivalent  in  grammes. 


HEAT  OF  CHEMICAL  ACTION.  181 

orimeter.  When  both  solutions  have  cooled  to 
the  same  temperature,  pour  them  together  into 
the  calorimeter  and  observe  the  rise  of  tem- 
perature. Calculate  for  one  molecule  HaS04  in 
grammes. 

Ex.  81.  Heat  of  Chemical  Action.—  Mix  250 
cubic  centimetres  of  water  with  fifty  cubic  centi- 
metres of  strong  sulphuric  acid,  and  when  the 
mixture  is  cold  place  it  in  the  calorimeter.  Scrupu- 
lously clean  a  strip  of  sheet  zinc  about  two  inches 
wide  by  five  inches  long.  Accurately  weigh  the 
zinc.  Plunge  the  strip  into  the  acid  and  allow 
the  action  to  continue  until  the  temperature  has 
risen  three  or  four  degrees.  Then  remove  the 
metal,  note  the  rise  of  temperature,  and  after 
washing  strip  with  water  and  alcohol,  dry,  and  de- 
termine the  loss  of  weight.  Calculate  the  amount 
of  heat  evolved  for  each  molecule  in  grammes  of 
ZnS04  formed. 

Ex.  82.  Heat  of  Precipitation.  —  Dissolve  a 
weighed  amount — about  twenty  grammes — of  crys- 
tallized baric  chloride  (BaCla .  2  H3O)  in  250  cubic 
centimetres  of  water.  Calculate  the  quantity  of 
sulphuric  acid  of  known  specific  gravity  required 
to  decompose  the  salt,  and  mix  the  acid  in  slight 
excess  of  the  calculated  amount  with  250  cubic 
centimetres  of  water.  Handle  the  solutions  as  in 
Ex.  80,  pouring  first  the  acid  and  afterwards  the 
salt  solution,  slowly  and  with  constant  stirring, 


182  LABORATORY  PRACTICE. 

into  the  calorimeter.  Calculate  for  each  molecule 
of  BaSO4  (in  gramme  units)  formed. 

Ex.  83.  Crystallized  Cupric  Sulphate. — Dis- 
solve a  weighed  amount — about  twenty  grammes 
—of  blue  vitriol  in  about  250  cubic  centimetres  of 
water  and  place  the  solution  in  the  calorimeter. 
Plunge  in  the  solution  a  strip  of  zinc  prepared  as 
in  Ex.  81  and  stir  until  the  copper  is  wholly  pre- 
cipitated, and  then  note  the  rise  of  temperature. 
Calculate  for  every  63 '6  or  every  atom  in  grammes 
of  copper  reduced. 

Unfortunately  the  fundamental  experiments 
on  this  subject — such,  for  example,  as  those  on 
the  heats  of  combustion  of  hydrogen,  carbon,  and 
sulphur,  or  the  heat  of  formation  of  hydrochloric 
acid — are  out  of  the  reach  of  elementary  students, 
and  indeed  of  most  teachers,  but  the  general  prin- 
ciples involved  can  be  readily  made  clear.  The 
chief  points  to  be  insisted  on  are : 

(1)  It  follows  from  the  principle  of  conservation  of  en- 
ergy, and  has  been  fully  proved  by  investigation,  that  in  a 
series  of  chemical  changes  the  total  amount  of  heat  devel- 
oped depends  wholly  on  the  initial  and  final  states  of  the 
system,  and  is  not  dependent  on  the  intermediate  steps. 
Thus  ninety-six  grammes  of  sulphuric  acid  consists  of  thirty- 
two  grammes  of  sulphur,  sixty-four  grammes  of  oxygen, 
and  two  grammes  of  hydrogen,  and  can  be  prepared  by  com- 
bining these  relative  amounts  of  roll  brimstone,  oxygen  gas, 
and  hydrogen  gas  in  several  different  ways.  But  whatever 
may  be  the  series  of  processes  employed,  the  total  amount  of 
heat  evolved  in  the  production  of  ninety-eight  grammes  of 
this  definite  compound  from  the  several  elementary  sub- 


EXOTHERMOUS  AND  ENDOTHERMOUS.    183 

stances  will  be  193,100  units.  Hence,  conversely,  if  by  a  se- 
ries of  analytical  processes  we  resolve  back  this  compound 
into  the  same  elementary  substances  in  the  same  condition 
an  equal  amount  of  heat  will  be  absorbed.  It  is  only  excep- 
tionally the  case  that  we  can  prepare  compounds  by  the 
direct  union  of  elementary  substances,  and  even  when  we 
can  it  is  rarely  that  we  can  measure  the  heat  thus  devel- 
oped ;  but  the  above  principle  renders  this  unnecessary. 
We  can  always  determine  the  heat  evolved  in  the  production 
of  a  compound  from  elementary  substances  (or  conversely) 
by  measuring  the  heat  evolved  at  each  step  of  the  successive 
operations  by  which  it  may  be  made,  and  we  are  thus  able 
to  choose  such  processes  as  are  adapted  to  thermal  measure- 
ments ;  and  in  the  investigations  of  thermo  -  chemistry  a 
great  deal  of  ingenuity  has  been  shown  in  tbis  selection  or 
in  devising  new  processes  which  are  compatible  with  the 
methods  of  calorimetry.  In  this  manner  what  is  termed 
the  heat  of  formation  of  a  large  part  of  known  compounds 
has  been  measured.  In  making  our  calculations  and  in 
stating  results  we  adopt  the  system  already  referred  to 
(Ex.  79),  and  a  table  giving  the  more  important  data 
will  be  found  in  the  author's  work  on  Chemical  Philos- 
ophy. 

The  thermal  relations  have  led  to  the  division  of  com- 
pounds into  two  large  classes — exothermous  bodies  (by  far 
the  larger  class),  whose  formation  from  known  elementary 
substances  in  their  familiar  state  is  attended  with  the  evo- 
lution of  heat,  and  endothermous  compounds,  of  which  the 
reverse  is  true. 

It  also  follows  from  the  principle  we  have  been  consid- 
ering that  if  we  begin  with  the  same  substance  in  the  same 
state  and  by  different  processes  reach  two  different  products, 
the  difference  in  the  heat  evolved  in  the  two  cases  is  that 
required  to  pass  from  one  product  to  the  other,  or,  what  is  an 
obvious  corollary,  if  the  initial  states  are  different  and  the 
final  results  in  all  respects  the  same,  then  the  same  relation 
will  hold  between  the  initial  states.  This  deduction  gives  us 
a  very  simple  means  of  determining  the  heat  of  combina- 
tion in  a  great  number  of  cases  where  direct  union  is  impos- 


184  LABORATORY  PRACTICE. 

sible  or  where  the  action  is  so  violent  that  all  thermal  meas- 
urements are  impracticable. 

Thus  28  grammes  of  olefiant  gas,  CaH4,  contain  24 
grammes  of  carbon  and  4  grammes  of  hydrogen,  and,  al- 
though we  can  not  combine  directly  charcoal  and  hydrogen 
gas,  we  can  determine  the  heat  evolved  in  the  production  of 
the  compound  in  this  way.  If  we  burn  28  grammes  of  olefi- 
ant gas  the  products  will  be  88  grammes  of  carbonic  diox- 
ide and  36  grammes  of  water — 

28  88  36 

CaH4  +  3O.  =  2CO2  +  2H2O  =  332,024. 

If  we  burn  24  grammes  of  charcoal  and  4  grammes  of  hydro- 
gen separately  we  shall  obtain  the  same  weights  of  the  same 
products  in  the  same  condition— 

24  88 

20  +  20,  =  2CO.  =  193,920 

4  36 

2H,  +  0,  =  2H,O  =  137,848 

331,768 

The  heat  evolved  in  all  these  three  processes  of  combus- 
tion has  been  measured  and  is  given  after  the  reaction.  The 
total  heat  evolved  in  burning  the  elementary  substances  is 
less  than  that  set  free  in  burning  the  compound  by  256  units. 
Hence  in  passing  from  charcoal  and  hydrogen  gas  to  olefi- 
ant gas  this  small  amount  of  heat  must  have  been  absorbed. 
Olefiant  gas  is  therefore  an  endothermous  compound. 

Again,  sulphuric  oxide  and  water  when  united  in  the 
proportions  indicated  by  the  reaction — 

SOs  +  HaO  =  HaSO4 

combine  with  explosive  violence,  but  we  can  readily  dis- 
solve both  SOs  and  HaSO4  in  an  equally  large  volume  of 
water  and  determine  the  heat  evolved  in  each  case.  The 
final  result  in  both  cases  is  a  weak  solution  of  sulphuric 
acid,  and  the  difference  between  the  amounts  of  heat  evolved 
gives  the  amount  which  would  be  given  by  the  above  reac- 
tion if  it  could  be  measured. 


TENDENCY  OF  CHEMICAL  PROCESSES.         185 

(2)  With  a  table  giving  the  heats  of  formation  of  the 
more  important  compounds  in  different  conditions  (whether 
solid,  liquid,  aeriform,  or  in  solution  in  water)  we  are  in  po- 
sition to  calculate  the  heat  evolved  in  any  ordinary  chemi- 
cal process.     We  have  only  to  compare  the  sum  of  the  heats 
of  formation  of  the  factors  of  the  reaction  with  that  of  the 
products,  paying  careful  regard  to  the  conditions  in  which 
the  several  materials  are  present.     Thus,  in  the  reaction  we 
have  discussed  so  often — 

(HaSO*  +  Aq)  +  Zn  =  (ZnSO*  +  Aq)  +  Hs, 

the  heat  evolved  during  the  process  is  the  difference  between 
the  heat  of  formation  of  sulphuric  acid  in  aqueous  solution 
and  that  of  zinc  sulphate  dissolved  in  an  equal  amount  of 
water.  We  need  pay  no  regard  to  the  elementary  sub- 
stances, either  the  zinc  dissolved  or  the  hydrogen  gas  set 
free,  for  they  are  present  in  the  very  condition  which  our 
calculations  assume,  and  since  the  heat  of  formation  of 
dilute  sulphuric  acid  is  210,000  and  that  of  the  solution  of 
zinc  sulphate  252,000  there  must  be  set  free  in  the  reaction — 

252,000  -  210,000  =  42,000  units. 

(3)  It  has  been  inferred  as  a  generalization  from  a  great 
number  of  facts  that,  other  things  being  equal,  the  activity 
of  a  chemical  process  is  proportional  to  the  amount  of  heat 
evolved,  and  that  where  several  courses  are  possible  the  tend 
ency  is  always  to  form  those  products  which  involve  the 
greatest  evolution  of  heat.     Often  by  restraining  the  reac- 
tion (as  by  lowering  the  temperature,  diluting  the  solution, 
or  restricting  the  amount  of  material)  other  products  may 
result ;  but  if  we  give  the  chemical  action  full  play  the  tend- 
ency is  as  above  stated.     In  the  action  of  nitric  acid  on  cop- 
per there  may  be  formed  either  nitric  oxide  or  nitrogen  gas, 
thus  : 

3Cu  +  (8HNO,  4-  Aq)  =  (3Cu(N03)a  +  4H,O  +  Aq)  +  2NO, 
or 

5Cu  +  (12HNO,  +  Aq)  =  (5Cu(NOs)3  -I-  6H3O  +  Aq)  +  N,. 


186  LABORATORY  PRACTICE. 

The  last  develops  the  most  heat,  and  nitrogen  gas  is  the  chief 
or  sole  product  formed  if  the  materials  are  allowed  to  hecome 
heated  ;  but  if  the  flask  is  kept  cool  the  chief  or  only  product 
is  nitric  oxide,  as  in  Ex.  44  (a). 

If  no  heat  would  be  set  free  by  an  assumed  process  the 
reaction  can  not  take  place  without  some  aid.  There  ap- 
pears no  reason  in  the  form  of  the  reaction  why  copper 
should  not  act  on  dilute  sulphuric  acid  like  zinc,  that  is  — 

Cu  +  (H3S04  +  Aq)  yield  (CuSO*  +  Aq)  +  H, 

But  while  the  heat  of  formation  of  an  aqueous  solution  of 
sulphuric  acid  is  210,000,  as  above,  that  of  an  aqueous  solu- 
tion of  copper  sulphate  is  199,100  (instead  of  252,000,  as  in 
the  case  of  zinc  sulphate),  and  heat  would  be  absorbed,  not 
evolved,  by  the  chemical  change. 

The  aid  required  to  determine  a  reaction  in  such  cases 
may  be  furnished  either  by  external  energy,  as  the  sun's 
rays  acting  on  the  green  foliage  of  the  vegetable  kingdom 
or  by  some  simultaneous  exothermous  process  which  en- 
trains the  other  and  supplies  the  necessary  heat.  Thus,  in 
the  above  reaction,  if  a  small  amount  of  nitric  acid  is  added 
(as  shown  in  Ex.  54  (&))  the  copper  at  once  dissolves  simply 
because  the  nitric  acid,  by  oxidizing  the  hydrogen  evolved 
(compare  Ex.  43  (6)),  generates  the  heat  required  to  render 
the  process,  as  a  whole,  exothermous. 

Thus  it  is  that  the  formation  of  endothermous  com- 
pounds becomes  possible.  Nitrous  oxide,  N2O,  is  such  a  com- 
pound. In  its  production  from  nitrogen  and  oxygen  gases 
18,000  units  of  heat  are  absorbed.  Its  tendency,  therefore 
(in  itself  alone),  is  to  fall  back  into  the  constituent  gases, 
when  the  same  amount  of  heat  would  be  evolved.  In  the 
reaction  by  which  nitrous  oxide  is  made  — 


80,700        118,800  -  18,000, 

the  heats  of  formation  are  printed  under  the  symbol,  and  it 
will  be  seen  that,  as  a  whole,  the  process  develops  100,800 
units  of  heat.  But,  as  will  also  be  noticed,  this  heat  wholly 
comes  from  the  oxidation  of  hydrogen  to  form  water,  which 


CAUSE  OF  INSTABILITY.  187 

is  sufficient  to  furnish  all  that  is  required  for  the  production 
of  N2O  and  still  have  a  large  excess  over  what  is  required 
to  determine  the  reaction.  Indeed,  unless  restrained  (by. 
keeping  the  temperature  at  the  lowest  possible  point),  this  re- 
action will  take  the  form — 

NH«NOs  =  2HaO  +  N,  +  iOa, 

which  corresponds  to  a  larger  evolution  of  heat,  and  to  just 
as  much  more  as  was  used  above  in  the  production  of  NaO. 

(4)  Endothermous  compounds  are  always  in  a  condition 
of  unstable  equilibrium,  and  sometimes  highly  explosive. 
This  is  strikingly  true  of  iodide  of  nitrogen,  which  often  ex- 
plodes at  the  mere  touch  of  a  feather  and  is  resolved  wholly 
into  elementary  substances — 

2NL  =  Na  +  3Ia. 

They  may  endure,  often  for  a  long  time,  in  consequence 
probably  of  features  of  molecular  structure  such  as  we  en- 
deavored to  illustrate  in  the  case  of  ammonium  nitrate,  but 
sooner  or  later  they  fall  into  more  stable  conditions.  They 
may  be  compared  to  a  vaulted  cathedral  roof  of  which  the 
stones  are  firmly  locked  together  and  held  high  in  air  by 
buttresses,  but  when  keystone  or  buttress  fail  fall  in  ruin. 

There  are  many  substances  which,  although  exother- 
mous  to  the  elementary  substances  to  which  their  heat  of 
formation  is  referred,  are  endothermous  in  their  relations  to 
certain  definite  products  into  which  they  are  more  or  less 
readily  resolved  with  evolution  of  heat,  or  in  their  rela- 
tions to  associated  material,  from  uniting  with  which  they 
are  restrained  by  physical  disabilities  or  conditions  of  struct- 
ure. Nitro-glycerin  and  gun-cotton,  already  referred  to,  are 
examples  of  the  first  type,  while  gunpowder,  in  which  com- 
bustible charcoal  and  sulphur  is  kept  apart  from  the  great 
store  of  oxygen  in  the  grains  of  nitre  by  the  inertness  of  the 
solid  state,  is  an  equally  striking  example  of  the  second  type. 
All  these  explosive  agents  owe  their  efficiency  not  only  to 
the  heat  evolved  by  the  internal  combustion,  which  ensues 
when  they  are  fired,  but  also  to  the  circumstance  that  the 
products  into  which  they  fall  are  for  the  most  part  aeriform 


188  LABORATORY  PRACTICE. 

bodies  whose  molecules  acquire  an  enormous  moving  power 
under  the  influence  of  the  heat  thus  generated. 

Of  instability  arising  from  association  by  far  the  most 
wonderful  example  is  furnished  by  the  presence  on  the  sur- 
face of  the  globe  of  a  large  amount  of  combustible  material 
in  contact  with  the  oxygen  of  the  atmosphere.  Almost  the 
whole  of  this  material  is  made  up  either  of  the  organized 
structure  of  plants  and  animals  or  else  of  the  remains  of 
such  structures,  and  however  multifarious  the  substances  of 
which  this  organic  matter  may  consist,  the  ultimate  ele- 
ments, with  unimportant  exceptions,  are  carbon,  hydrogen, 
nitrogen,  and  oxygen,  and  the  whole  of  this  material  was 
primarily  formed  from  the  constituents  of  air  and  water, 
including,  of  course,  carbonic  acid  and  ammonia,  always 
present  in  the  atmosphere  and  in  the  water  permeating  the 
soil.  From  these  materials  the  plant  obtains  almost  its  sole 
food,  and  the  animal  ultimately  at  least  lives  on  the  plant. 
In  some  mysterious  way  the  sun's  rays  give  the  plant  the 
power  of  producing  the  substances  of  its  tissues  from  the 
simple  articles  of  its  diet.  Of  the  manner  in  which  the 
highly  complex  products  are  built  up  we  know  almost  noth- 
ing. But  of  this  we  are  sure.  The  energy  of  the  sun's  rays 
is  the  power  by  which  carbonic  acid  and  water  are  decom- 
posed and  materials  so  unstable  in  the  presence  of  the  atmos- 
phere constructed.  Probably  the  effect  is  indeed  a  result  of 
molecular  construction  and  the  structure  endures  until  a 
conflagration,  or  the  slower  processes  of  decay,  destroys  the 
fabric  and  resolves  the  organic  matter  into  the  elements 
from  which  it  sprung.  Thus  there  is  a  constant  cycle  in 
nature.  The  sun's  rays  are  ever  building  up,  and  in  so  do- 
ing are  setting  free  the  very  oxygen  which,  sooner  or  later, 
will  destroy  all  this  work.  We  also  know  that  the  heat 
given  out  in  burning  is  the  exact  equivalent  of  the  work 
done  in  building,  and  therefore  is  simply  transmuted  solar 
energy.  Just  as  the  sun  lifts  the  water  whose  fall  main- 
tains the  great  aqueous  circulation  of  the  globe,  so  the  same 
vitalizing  energy  builds  up  organic  structures,  by  whose 
reabsorption  into  the  all-devouring  atmosphere  the  life  and 
activity  on  the  planet  is  sustained.  Moreover,  only  by  util- 


EXPLOSIVES.  189 

izing  this  same  energy— though  often  so  indirectly  that  it  es- 
capes notice — are  we  able  to  produce  endothermous  or  un- 
stable compounds  in  our  laboratories. 

(5)  It  must  be  remembered  that  compounds  are  endoth- 
ermous only  in  relation  to  the  elementary  substances  in  a 
definite  condition  (carbon  as  charcoal  sulphur,  as  brimstone, 
etc.),  from  which  in  our  system  they  are  regarded  as  having 
been  formed.  Their  peculiar  thermal  relations  do  not  neces- 
sarily imply  a  want  of  chemical  energy  between  the  element- 
ary atoms  of  which  their  molecules  are  supposed  to  consist. 
Thus  the  oxides  of  nitrogen  are  endothermous,  and  yet  there 
can  be  no  question  that  the  nitrogen  atoms  have  a  marked 
affinity  for  the  atoms  of  oxygen.  If  we  could  deal  with  ele- 
mentary atoms  all  compounds  would  be,  doubtless,  exother- 
mous,  but  when  we  deal  with  the  elementary  substances  in 
their  normal  condition  the  constitution  of  the  molecules  of 
these  elementary  substances  comes  into  play.  As  has  been 
shown,  the  molecules  of  elementary  substances,  as  well  as 
those  of  compounds,  are  aggregates  of  atoms,  only  of  atoms 
all  of  which  are  of  the  same  kind,  and  not,  as  in  compound 
substances,  of  different  kinds.  The  molecule  of  oxygen  gas 
(Oa)  is  formed  of  two  atoms  of  oxygen  ;  that  of  ozone  (Os), 
of  three  atoms  of  oxygen  ;  that  of  nitrogen  gas  (N2),  of  two 
atoms  of  nitrogen,  etc.  The  atoms  of  nitrogen,  united  in  a 
molecule  of  the  gas  must  have  an  attraction  for  each  other, 
else  they  would  not  so  group  themselves  in  pairs.  As  yet 
we  have  not  been  able  to  measure  this  attraction  with  confi- 
dence ;  but  we  have  good  reason  for  believing  that  it  is  very 
strong.  The  endothermous  compounds  called  nitrogen 
iodide  (NI8)  and  nitrogen  chloride  (NCls)  are  so  explosive 
not,  as  we  believe,  because  the  nitrogen  atoms  have  no  affin- 
ity for  the  iodine  and  chlorine  atoms,  but  because  they  have 
such  a  strong  attraction  for  each  other  that  they  break  from 
the  iodine  and  chlorine  atoms  and  rush  together. 

2NL  =  N.  +  31, 
2NC1,  =  N,  +  3C1, 

We  explain  the  inertness  of  nitrogen  gas  in  this  way. 
Simply  the    nitrogen  atoms   exert   a  <  stronger    attraction 
13 


190  LABORATORY  PRACTICE. 

among  themselves  than  for  those  of  the  other  chemical  ele- 
ments with  only  a  few  exceptions.  Hence,  also,  the  general 
instability  of  the  compounds  of  nitrogen  as  a  class.  Animal 
structures  are  for  the  most  part  made  up  of  nitrogenized 
substance,  and  thus  decay  and  death  in  nature  are  closely 
associated  with  this  striking  feature  of  our  chemical  phi- 
losophy. 

(6)  Could  we  experiment  with  isolated  atoms  all  chemi- 
cal relations  would   unquestionably   appear   simpler,  and 
modern  science  has  rendered  probable  that    there    exists 
such  a  condition  in  the  universe.     It  is  well  known  that 
heat  tends  to  decompose  chemical  compounds,  and  the  phe- 
nomena thus  resulting  form  a  very  interesting  subject  of 
chemical  inquiry  known   as    thermolysis   or   dissociation. 
Steam  passed  through  metal  tubes  at  a  white  heat  acts  in 
every  respect  like  a  mixture  of  oxygen  and  hydrogen  gases, 
and,  as  common  experience  shows,  most  compounds  even 
at  a  red  heat  suffer  more  or  less  fundamental  chemical 
changes.     The  known  facts  point  to  the  conclusion  that  at 
such  high  temperatures  as  must  rule  at  the  sun  and  at  the 
fixed  stars  all  known  materials  would  be  resolved  into  ele- 
mentary atoms.      Spectroscopic  observations  confirm  this 
inference,  and  such  evidence  has  even  been  interpreted  as 
indicating  that  some  of  the  atoms  which  we  regard  as  ele- 
mentary are  resolved  into  still  simpler  parts  at  the  great 
focus  of  solar  radiation.     If  the  nebular  hypothesis  is  cor- 
rect our  world  must  have  been  primarily  in  this  condition, 
and  the  substances  which  we  now  find  on  its  crust  must 
have  been  formed  as  the  elementary  atoms  came  together  in 
the  process  of  cooling  and  united  in  accordance  with  their 
mutual  affinities.    According  to  this  hypothesis,  the  original 
chaos  out  of  which  the  present  order  sprang  was  a  condition 
of  isolated  atoms,  and  the  foundations  of  the  globe  must 
have  been  laid  in  flames. 

(7)  In  this  last  chapter  of  our  book  we  have  only  been 
able  to  illustrate  a  few  of  the  more  important  relations  of 
thermo-chemistry.     Our  one  object  has  been  to  exhibit  the 
scope  of  this  department  of  chemical  science,  as  we  have 
sought  to  show  in  earlier  chapters  that  of  qualitative  and 


CONCLUSIONS.  191 

of  quantitative  analysis.  The  theoretical  relations  of  the 
science  have  been  previously  set  forth  by  us  in  popular 
form  in  the  New  Chemistry.*  It  is  expected  that  this  will 
serve  as  a  companion  to  the  former  book,  and  the  student 
who  thoughtfully  performs  all  the  experiments  and  dili- 
gently inquires  what  each  is  calculated  to  teach  can  not  fail 
to  gain  clear  ideas  of  the  methods  of  chemical  investigation 
and  at  the  same  time  will  acquire  skill  in  drawing  inferences 
from  experimental  data.  Having  thus  seen  the  relations  of 
the  broader  divisions  of  the  subject,  the  student  will  be  pre- 
pared to  enter  on  the  professional  study  of  chemistry  intelli- 
gently, or,  if  he  goes  no  further,  will  have  acquired  a  clear 
conception  of  the  aims  and  methods  of  the  science  and  of  its 
true  position  in  a  scheme  of  education.  The  next  step  in 
the  study  of  chemistry  should  be  to  acquire  an  adequate 
knowledge  of  the  scheme  of  the  chemical  elements  as  it  is 
presented  by  Roscoe  and  Schorlemmer,f  or,  still  better,  as  it 
is  illustrated  experimentally  in  an  extended  course  of  lect- 
ures such  as  is  given  every  year  at  Harvard  or  at  any  one  of 
our  principal  universities.  This  is  a  serious  task,  since  the 
mass  of  details  is  very  great  and  the  subject  can  not  be 
profitably  abridged  beyond  a  limited  extent.  A  brief  epit- 
ome will  not  be  of  much  value.  The  mind  must  dwell  on 
the  subject  in  its  various  relations  in  order  to  make  the 
knowledge  real  or  lasting,  and  unless  the  student  has  that 
object  in  the  acquisition  which  will  lead  him  to  give  to  the 
work  the  requisite  time,  he  had  better,  especially  if  it  is  a 
question  of  liberal  education,  limit  his  study  to  the  general 
principles  of  the  science  as  presented  in  this  or  some  similar 
book. 

*  D.  Appleton  &   Co.,  publishers.     New  York.     Last  edition, 
1890. 

f  Treatise  on  Chemistry,  D.  Appleton  &  Co. 


192 


LABORATORY  PRACTICE. 


List  of  the  Elementary  Substances,  excepting  such  as  are 
very  Rare  or  of  Doubtful  Authenticity. 


Aluminum, 

Al,    . 

...     27-1 

Molybdenum, 

Mo    .. 

96 

Antimony, 

Sb, 

...  120 

Nickel, 

Ni 

39 

Arsenic 

75 

Nitrogen 

N 

14 

Barium, 

Ba, 

...  137 

Osmium, 

Os 

190*9 

Bismuth, 

Bi, 

...208 

Oxvsren 

o 

16 

Boron 

B 

...     11 

Palladium 

Pd 

106  '6 

Bromine, 
Cadmium, 

Br,    . 
Cd,    .. 

...     80 
...  112-2 

Phosphorus, 
Platinum 

*•  **i   •  •  •  • 
P,     .... 
Pt 

31 
194-8 

Caesium 

Cs 

.  .  .  133 

Potassium 

K 

39-1 

Calcium 

Ca 

40 

Rhodium 

Rh 

103 

Carbon, 

c, 

...     12 

Rubidium 

Rb 

85  '4 

Cerium 

Ce 

.  .  .  140 

Ruthenium 

Ru 

101  -8 

Chlorine 

Cl 

35-5 

Scandium 

J-VLt,       .... 

Sc 

44 

Chromium, 

Cr, 

...     52-1 

Selenium 

KJV,    .... 
Se 

79*2 

Cobalt 

Co 

...     59 

Silicon 

Si 

28-3 

Columbium 

Cb 

.     94 

Silver 

As- 

108 

Copper. 

Cu,    . 

...     63-6 

Sodium 

**•!     .... 

Na    . 

23 

Didymium,  I 
Erbium 

*d,  141 
Er 

Pr,  144 
166 

Strontium, 
Sulphur 

Sr,    .... 
S 

87-6 
32*1 

Fluorine, 

F, 

...     19 

Tantalum 

uj      .... 
Ta 

182 

Gallium 

Ga 

...     70 

Tellurium 

Te 

125  ? 

Germanium 

Ge 

72'3 

Terbium  ? 

Tr 

171  ? 

Glucinum, 

Gl, 

9-1 

Thallium 

j-i,    .... 
Tl 

204  •  1 

Gold 

Au 

...  197*2 

Thorium 

Th 

232 

Hydrogen 

H 

1 

Thulium  ? 

Tm 

171 

Indium, 
Iodine 

In,     .. 
I 

...  113-7 
.  .  .  126  •  9 

Tin, 
Titanium 

Sn,    .... 
Ti 

118 

48 

Iridium 

Ir      . 

193 

Tungsten 

W 

184 

Iron 

Fe 

56 

TJran  ium 

Ur 

240 

Lanthanum 

La 

139 

"Vanadiu  m 

w*i 

Va 

51'4 

Lead, 

Pb,    .. 

...  206-9 

Ytterbium  ? 

v  d,      .... 

Yb 

173 

Lithium 

Li 

7 

Yttrium 

Y 

90 

Magnesium 

Ms- 

24*4 

Zinc 

Zn 

65*2 

Manganese, 
Mercurv. 

•"•*-&  i  • 
Mn,  . 
He-.   . 

...     55 
.  200 

Zirconium, 

Zr,     .... 

90 

THE   END. 


KEROSENE  STOVE  AND  TUBE  FURNACE. 

Well  adapted  for  most  cnemical  experiments  referred  to  in  this  book  (Ex. 


(1) 


TWENTIETH  CENTURY  TEXT-BOOKS/ 

Elementary  Chemistry. 

By  ROBERT  HART  BRADBURY,  A.M.,  Ph.D., 
Teacher  of  Chemistry  in  the  Philadelphia  Central 
Manual  Training- School ;  formerly  Lecturer  on 
Physical  Chemistry  in  the  Department  of  Phi- 
losophy, University  of  Pennsylvania.  I2mo.  Cloth, 
$1.00.  With  Laboratory  Manual,  $1.25. 

There  is  no  dearth  of  text-books  of  Chemistry,  but  all  seem  to  have 
fallen  short  in  some  essential  particulars,  and  fail  to  meet  the  actual 
requirements  of  class  work.  This  book  attempts  to  remedy  such  faults. 
It  is,  therefore,  thoroughly  fresh  in  material  and  treatment.  It  is  not  a 
compilation  from  existing  manuals,  but  the  result  of  a  survey  of  the 
recent  literature  of  the  science  ;  yet  nothing  has  been  introduced  that 
has  not  been  tested  practically  in  the  class-room  and  found  to  be  avail- 
able. Special  effort  has  been  made  to  attain  precision  and  clearness  of 
statement.  The  theoretical  and  general  portions  of  the  subject  are  not 
hurled  at  the  student  in  an  overwhelming  and  unintelligible  mass.  The 
introduction  of  each  generalization  is  deferred  until  a  particular  case  of 
it  has  been  encountered.  Formulae  and  equations  are  not  introduced 
until  the  student  is  able  to  receive  them,  and  the  atomic  theory  is  re- 
served until  the  student  is  able  to  understand  it. 

NOTE. — This  text-book  was  scarcely  off  the  press  when  it  was  adopted  in 
Boys'  High  School,  New  York  (the  largest  high  school  in  the  world). 
High  School  of  Commerce,  New  York. 
Erasmus  Hall  High  School,  Brooklyn. 
Manual  Training  High  School,  Brooklyn. 
Central  Manual  Training-School,  Philadelphia. 
High  School,  Hartford,  Conn. 

A  Laboratory  Manual  of  Chemistry. 

By  ROBERT  HART  BRADBURY,  A.M.,  Ph.D. 
Cloth,  45  cents. 

The  laboratory  guide  aims  to  supply  an  outline  of  work  which  can 
be  completed  by  the  average  student  in  one  school  year.  Constant 
questions,  problems,  and  exercises  serve  to  prevent  the  student's  work 
from  becoming  thoughtless  or  mechanical. 

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NEW  YORK.      BOSTON.     CHICAGO.     LONDON. 


TWENTIETH   CENTURY   TEXT-BOOKS. 

PHYSICS. 

Elements  of  Physics. 

By  C.  HAN  FORD  HENDERSON,  A.  M.,  Ph.  D.,  and  JOHN 
F.  WOODHULL,  Ph.  D.,  Professor  of  Physics  in  Teachers 
College,  Columbia  University.  i2mo.  Cloth,  $1.10.  With 
Experiments,  $1.25. 

Physical  Experiments :    A  Laboratory  Manual. 

By  JOHN  F.  WOODHULL,  and  M.  B.  VAN  ARSDALE, 
Instructor  in  Physical  Science  in  Horace  Mann  School  and 
Assistant  in  Teachers  College.  i2mo.  Cloth,  45  cents; 
with  alternate  blank  pages,  60  cents. 

Inexpensive  Apparatus.  The  figures  used  to  illustrate  these 
books  and  the  apparatus  recommended  in  them  are  comparable  to 
maps  that  show  only  the  matter  under  consideration  and  leave  out  the 
things  that  confuse,  divert,  and  obscure. 

Self-Explanatory.  The  text-book  is  not  a  syllabus  of  rules  as 
an  adjunct  to  laboratory  experiments,  nor  merely  a  table  of  reference, 
but  is  self-explanatory  and  "  readable  "  in  the  best  sense,  its  literary 
style  being  lucid,  fresh,  warm,  and  attractive.  Hence  it  gives  the  sub- 
ject a  live  interest. 

Laboratory  Work.  The  laboratory  manual  is  something  more 
than  a  book  of  directions  for  laboratory  work.  It  is  full  of  suggestions 
of  facts  and  phenomena  in  nature  parallel  with  the  experiments. 
Problems  are  very  abundant,  drawn  largely  from  sources  within  the 
pupil's  own  experience,  and  intended  to  lead  the  pupil  to  organize  his 
thoughts  in  connection  with  his  laboratory  work. 

Special  Chapters.  The  treatment  of  electricity  in  these  books 
ij  particularly  clear  and  simple.  It  treats  of  the  telegraph,  the  electric 
bell,  the  electric  light,  the  motor,  the  dynamo,  electric  cars,  telephone, 
wireless  telegraphy,  Roentgen  rays,  etc.  The  chapters  on  Mechanics, 
Heat,  Light,  and  Sound  also  areremarkably  interesting  as  well  as 
highly  instructive. 

Portraits  of  eminent  scientists  and  brief  statements  concerning 
their  lives  and  labors  make  the  study  more  attractive  by  giving  it  £. 
concrete  human  aspect. 

The  text-book  and  laboratory  manual  meet  perfectly  ever*? 
detail  of  college-entrance  requirements. 

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LABORATORY   MANUAL   OF   PHYSICS. 

Laboratory    Exercises    in    Elementary 
Physics. 

A  Manual  for  Students  in  Academies  and 
High  Schools.  By  FRANKLIN  H.  AYRES,  In- 
structor in  Physics  in  Central  High  School, 
Kansas  City,  Mo.  i2mo.  Cloth,  60  cents. 


SIX  PARTICULARS. 

1.  This  Manual  is  strictly  in  line  with  the  best  university 

practise. 

2.  It  keeps  within  the  limits  of  legitimate  laboratory  work 

in  elementary  physics. 

3.  Its  exercises  demand  work  and  thought  worthy  of  the 

best  student,  but  within  the  range  of  the  average. 

4.  It  is  perfectly  clear  with  respect  to  (a)  The  general 

object  of  the  exercises  ;  (b)  How  the  work  is  to  be 
done  ;  and  (c)  The  road  the  student  must  travel  to 
arrive  at  the  conclusions  involved. 

5.  Its  course  in  measurements  is  founded  on  the  cardinal 

principle  that  no  physical  law  with  a  quantitative 
aspect  can  be  understood  until  the  beginner  can 
express  in  mathematical  form  both  the  law  itself  and 
the  process  by  which  he  verified  it. 

6.  It  presents  laboratory  physics  in  the  spirit  and  by  the 

practise  of  pure  induction,  yet  carefully  avoids  the 
absurdity  of  supposing  that  young  students  are  capa- 
ble of  original  investigation. 

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TWENTIETH    CENTURY    TEXT-BOOKS. 
A  Text-Book  of  Geology. 

By  Professor  ALBERT  PERRY  BRIGHAM,  of  Colgate  Uni- 
versity. 477  Pages.  295  Illustrations.  121110.  Cloth,  $1.40. 

This  superb  text-book  is  the  best  account  for  secondary 
schools  of  the  earth's  marvelous  origin,  of  the  processes  that 
brought  the  ordered  world  out  of  chaos,  and  of  the  phenomena 
of  geologic  evolution — considered  dynamically,  structurally,  and 
historically.  The  planet's  life  history  is  told  with  directness, 
brevity,  and  pedagogic  fitness.  The  text  is  supplemented  with 
295  exquisite  photographic  illustrations,  many  taken  by  Professor 
Brigham  for  this  work.  An  exceptional  success  in  text-book 
writing  and  text-book  making. 

"  Brigham's  Geology  is  the  cleanest  cut  and  best  pedagogical  text- 
book for  the  high  school  that  I  have  seen." — C.  H.  Richardson, 
Hanover,  N.  H. 

"  Most  interesting.  Decidedly  the  most  practical  book  that  I 
have  seen  for  use  in  high  schools." — Miss  S.  A.  Edwards,  Philadelphia 
High  School  for  Girls. 

I  consider  it  the  best  written  and  best  illustrated  book  I  have  ever 
seen  for  secondary  schools." — C.  F.  Warner,  Mechanics  Arts  High 
School,  Springfield,  Mass. 

"  The  most  attractive  text-book  of  Geology  for  secondary  schools 
that  I  have  seen.  The  illustrations  are  a  delight." — Belle  Sherman, 
Ithaca  High  School,  Ithaca,  N.  Y. 

"  It  is  magnificent.  I  consider  it  superior  to  any  other  book  of  the 
kind  in  illustrations,  text,  and  adaptation  to  field  work." — Mrs.  L.  L. 
W.  Wilson,  Philadelphia  Normal  School. 

"  In  every  way  fully  equal  to  any  of  the  splendid  series  of  Twen- 
tieth Century  Text-Books.  Many  of  the  illustrations  are  new  and 
their  execution  is  perfect."—^.  /.  Schiedt,  Professor  of  Geology, 
Franklin  and  Marshall  College,  Lancaster,  Pa. 

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TWENTIETH  CENTURY  TEXT-BOOKS. 


An  Introduction  to  Physical  Geography. 

By  GROVE  KARL  GILBERT,  LL.  D.,  United  States  Geological  Sur- 
vey ;  Author  of  "The  Geology  of  the  Henry  Mountains,"  "Lake 
Bonneville,"  Numerous  Reports,  etc.,  in  publications  of  United 
States  Geological  Survey ;  and  ALBERT  PERRY  BRIGHAM,  A.  M., 
Professor  of  Geology,  Colgate  University,  Hamilton,  N.  Y.,  Fellow 
of  the  Geological  Society  of  America,  etc.,  Associate  Editor  Bulletin 
American  Geographical  Society,  Author  of  "  A  Text-Book  of 
Geology "  (Twentieth  Century  Text-Books).  Illustrated.  I2mo. 
Cloth,  $1.25. 

SIX    SALIENT   POINTS. 

The  new  pedagogy  of  Physical  Geography  receives  in  this  book  its  first 
adequate  presentation. 

Hence,  this  text  meets  the  present  requirements  of  high  school  and 
college-entrance  work  perfectly  and  in  full  detail. 

Treatment  adapted  to  the  early  years  of  the  course — the  book  will  interest 
pupils  aged  fourteen. 

Statements  throughout  are  not  merely  theoretical,  but  definitely  concrete, 
appropriately  illustrated,  and  logically  summarized. 

Topics  cover  "The  Physical  Environment  of  Man:"  The  Earth  as  a 
Globe,  the  Ocean,  the  Air,  and  the  Land — in  increasing  proportion. 

The  exquisite  half-tone  illustrations  far  surpass  in  beauty,  helpfulness, 
and  number  anything  before  attempted.  A  most  important  and  significant 
feature. 

THE   IDEAL  COURSE   AND   GILBERT  AND 

BRIGHAM'S   BOOK. 

This  book  meets  fully,  in  minute  detail,  and  for  the  first  time,  all 
the  specifications  set  forth  in  the  Report  of  the  Committee  on  College 
Entrance  Requirements  to  the  National  Educational  Association  in 
1899.  It  keeps  accurately  to  the  definition  laid  down  ;  it  furnishes  the 
requisite  kind  and  amount  of  instruction  to  train  the  observation  and 
to  prepare  for  later  special  courses  in  science  ;  and  it  elevates  physical 
geography  beyond  cavil  to  the  proper  plane  for  a  college-entrance 
requirement,  by  organizing  its  content  to  its  highest  capacity  as  a 
pedagogic  discipline. 

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NEW  EDITION  OF  PROFESSOR  HUXLEVS  ESSAYS. 

Collected  Essays. 

By  THOMAS  H.  HUXLEY.  New  complete 
edition,  with  revisions,  the  Essays  being 
grouped  according  to  general  subject.  In 
nine  volumes,  a  new  Introduction  accom- 
panying each  volume.  I2mo.  Cloth,  $1.25 
per  volume. 

Vol. 

I.  Methods  and  Results. 
II.  Darwiniana. 

III.  Science  and  Education. 

IV.  Science  and  Hebrew  Tradition. 
V.  Science  and  Christian  Tradition. 

VI.  Hume. 

VII.  Man's  Place  in  Nature. 
VIII.  Discourses,  Biological  and  Geological. 
IX.  Evolution  and  Ethics,  and  Other  Essays. 

"  Mr.  Huxley  has  covered  a  vast  variety  of  topics  during 
the  last  quarter  of  a  century.  It  gives  one  an  agreeable  surprise 
to  look  over  the  tables  of  contents  and  note  the  immense  territory 
which  he  has  explored.  To  read  these  books  carefully  and 
itudiously  is  to  become  thoroughly  acquainted  with  the  most 
advanced  thought  on  a  large  number  of  topics." — New  Yqrk 
Herald. 

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THE 

INTERNATIONAL  SCIENTIFIC  SERIES. 


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CIERS. By  J.  TYNDALL.  LL.D..F.R.S.  With  35  Illustrations.  SI  .50. 

2  PHYSICS  AND  POLITICS;  or,  Thoughts  on  the  Application  of  the  Prin- 
ciples of  "Natural  Selection"  and  "Inheritance"  to  Political  Society. 
By  WALTER  BAGEHOT.  $1.50. 

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trations.    $1.75. 

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With  31  Illustrations.         $2.00. 

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PETTIGREW,  M.  D.,  F.  R.  S.,  etc.     With  130  Illustrations.     $1.75. 

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24.  A  HISTORY  OF  THE  GROWTH  OF  THE  STEAM-ENGINE.     By  Pro- 

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27.  THE  HUMAN  SPECIES.     By  Professor  A.  DE  QUATREFAGES,  Museum  of 

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28.  THE  CRAYFISH:  An  Introduction  to  the  Study  of  Zoology.     By  T.  H. 

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30.  ANIMAL  LIFE  AS  AFFECTED  BY  THE  NATURAL  CONDITIONS  OF 

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132  Illustrations.     $1 .50. 

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37.  THE  FORMATION  OF  VEGETABLE  MOULD,  THROUGH  THE  AC- 
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40.  MYTH  AND  SCIENCE.     By  TITO  VIGNOLI.     $1.50. 

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44.  ANIMAL  INTELLIGENCE.  By  GEORGE  J.  ROMANES,  M.  D.,  F.  R.  S.  $1.75. 

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60.  INTERNATIONAL  LAW,  with  Materials  for  a  Code  of  International  Law. 

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62.  ANTHROPOLOGY.    An  Introduction  to  the  Study  of  Man  and  Civilization. 

By  EDWARD  B.  TYLOR,  D.  C.  L.,  F.  R.  S.     With  78  Illustrations.     82.00. 

63.  THE  ORIGIN  OF  FLORAL  STRUCTURES,  THROUGH  INSECT  AND 

OTHER  AGENCIES.    By  the  Rev.  GEORGE  HENSLOW,  M.  A.,  etc.    With 

88  Illustrations.     $1.75. 

64.  THE    SENSES,    INSTINCTS,    AND    INTELLIGENCE    OF    ANIMALS, 

WITH  SPECIAL  REFERENCE  TO  INSECTS.     By  Sir  JOHN  LUBBOCK, 
Bart.,  F.  R.  S.,  etc.     With  118  Illustrations.     $1.75. 

65.  THE  PRIMITIVE  FAMILY  IN  ITS  ORIGIN  AND  DEVELOPMENT. 

By  Dr.  C.  N.  STARCKE,  University  of  Copenhagen.     $1.75. 

66.  PHYSIOLOGY  OF  BODILY  EXERCISE.    By  F.  LAGRANGE,  M.  D.    $1 .75. 

67.  THE  COLORS   OF   ANIMALS:    Their   Meaning   and   Use.     By   EDWARD 

BAGNALL  POULTON,  F.  R.  S.     With  36  Illustrations  and  1  Colored  Plate. 
$1.75. 

68.  SOCIALISM:    New   and    Old.     By    Professor    WILLIAM    GRAHAM,    M.  A., 

Queen's  College,  Belfast.     $1.75. 

69.  MAN    AND    THE    GLACIAL    PERIOD.     By    Professor    G.    FREDERICK 

WRIGHT,  D.  D.,  Oberlin  Theological  Seminary.     With  108  Illustrations 
and  3  Maps.     $1.75. 

71.  A    HISTORY    OF   CRUSTACEA.     Recent    Malacostraca.     By    the    Rev. 

THOMAS  R.  R.  STEBBING,  M.  A.     With  51  Illustrations.     $2.00. 

72.  RACE  AND  LANGUAGE.    By  Professor  ANDRE  LEFEVRE,  Anthropological 

School,  Paris.     $1.50. 

73.  MOVEMENT.     By  E.  J.  MAREY.     Translated  by  ERIC  PRITCHARD,  M.  A., 

M.B.,  B.Ch.  (Oxon.).     With  200  Illustrations.     $1.75. 

74.  ICE-WORK,  PRESENT  AND  PAST.     By  T.  G.  BONNEY,  D.  Sc.,  F.  R.  S., 

F.  S.  A.,  etc.,  Professor  of  Geology  at  University  College,  London.     $1.50. 

75.  WHAT  IS  ELECTRICITY?     By  JOHN  TROWBRIDGE,  S.  D.,  Rumford  Pro- 

fessor and  Lecturer  on  the  Applications  of  Science  to  the  Useful  Arts,  Har- 
vard University.     Illustrated.     $1 .50. 

76.  THE  EVOLUTION  OF  THE  ART  OF  MUSIC.     By  C.  HUBERT  H.  PARRY, 

D.  C.  L.,  M.  A.,  etc.     $1.75. 

77.  THE  AURORA  BOREALIS.     By  ALFRED  ANGOT,  Honorary  Meteorologist 

to  the  Central  Meteorological  Office  of  France.     $1.75. 

78.  MEMORY  AND  ITS  CULTIVATION.     By  F.  W.  EDRIDGE-GREEN,  M.  D., 

F.R.C.S.     $1.50. 

79.  EVOLUTION   BY  ATROPHY.     By  JEAN  DEMOOR,  JEAN  MASSART,  and 

EMILE  VANDERVELDE. 


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